SI units - set of 7 basic units from which all other units are derived  

• kilogram (kg) - measures mass (amount of material in an object)
• meter (m) - length
• second (sa) - time
• Kelvin (K) - temperature, where 0° is absolute zero; 273.15 more than Celsius
• mole (mol) - amount of a substance
• ampere (A) - electric current
• candela (cd) - luminosity

prefixes of metric system  

• Giga (G) - 109
• Mega (M) - 106
• Kilo (k) - 103
• Deci (d) - 10-1
• Centi (c) - 10-2
• Milli (m) - 10-3
• Micro (µa) - 10-6
• Nano (n) - 10-9
• Pico (p) - 10-12
• Femto (f) - 10-15

derived SI units - uses a combination of SI units  

• volume - measured by cubic meter or liter (equal to cubic decimeter)
• density - amount of mass per unit volume; mass/volume

matter - physical material of the universe  

• different compositions due to different arrangements of elements
• atoms - smallest building blocks of matter
• properties of matter change due to composition/structure of the atoms
• molecule - 2+ atoms joined in a specific shape

states of matter - gas, liquid, or solid  

• gas - aka vapor, has no definite shape/volume; molecules move at high speeds, often colliding w/ each other
• liquid - definite volume, but no definite shape; molecules can slide over each other
• solid - definite shape/volume, molecules held tightly together

pure substances - has distinct properties and doesn't vary from sample to sample  

• either elements or compounds
• elements - cannot be decomposed into simpler substances
  • oxygen, silicon, aluminum, iron, calcium make up most of crust
  • oxygen, carbon, hydrogen makes up most of human body
  • similar elements grouped in the same column on the periodic table
• compounds - combination of 2+ elements
  • law of constant composition (definite proportions) - composition of a pure compound is always proportionally the same
• mixture - combination of 2+ substances; each substance keeps its characteristics
  • components - substances making up a mixture
  • homogeneous mixture (solution) - same throughout
  • heterogeneous mixture - not the same composition throughout

properties of matter - can be physical or chemical  

• physical changes - can be measured w/o changing the composition of the substance (ex. odor, color, density, melting point, boiling point, hardness)
• chemical properties - describes how the substance reacts w/ other substacnes (ex. flammability)
• intensive property - doesn't depend on amount of substance available
• extensive property - depends on the amount of substance available
• physical change - substance retains composition but changes appearance
• chemical change - substance changes to a chemically different substance

separation of mixtures - possible by using the components' different properties  

• filtration - uses a filter to separate a mixture (usually water-solid mixture)
• substances w/ lower boiling points are more volatile
• distillation - boiling a solution to isolate the more volatile substance
• chromatography - "writing of colors"; depends on ability of substances to adhere to surfaces of various solids

scientific method - guidelines for the practice of science  

• starts w/ data collection
• hypothesis - temporary explanation for something; can be disproved
• scientific law - verbal/mathematical statement that summarizes lots of observations
• theory - general hypothesis that can consistently predict results

numerical uncertainty - measurements are ALWAYS inexact  

• precision - how closely measurements agree w/ each other
• accuracy - how closely measurements agree w/ the right answer
• precision doesn't lead to accuracy
• trials are done to get a better accuracy

significant figures - all digits of a quantity, including an uncertain one  

• measured quantities are reported so that the last digit is uncertain
• more significant figures means more certainty in the measurement
• all non-zero digits are significant
  • 3456 - 4 significant figures
  • 3.13245 - 6 significant figures
  • 22 - 2 significant figures
• zeros between non-zero digits are significant
  • 34000005 - 8 significant digits
  • 345.0003 - 7 significant digits
  • 60.7 - 3 significant digits
• zeros at the beginning of the number aren't significant
  • 0.00000000006 - 1 significant digit
  • 0000234 - 3 significan digits
  • 0.0005042 - 4 signifcant digits
• zeros at the end of the number are significant
  • 34000. - 5 significant digits
  • 34.00 - 4 significant digits
  • 5.01000 - 6 significant digits
• zeros at the end of numbers w/o a decimal point may or may not be significant
  • 450 - 2 or 3 significant digits
  • 3000 - 1, 2, 3, or 4 significant digits
• w/ multiplication/division, the result should have the same number of significant figures as the number w/ the least significant figures
• w/ addition/subtraction, result can't have more decimal places than the number w/ the least number of decimal places

dimensional analysis - guarantees that the result has the proper units  

• conversion factor - fraction where the numerator/denominator are the same quantity in different units
• given unit x desired unit / given unit = desired unit
• multiple conversion factors can also be used
• derived SI units may require multiple conversion factors to get the desired unit
• Convert 0.076L to mL
  • 0.076L x 1000mL / 1L = 76 mL
• Convert 1.55kg/m3 to g/L
  
  • 1.55kg / 1m3 x 1m3 / 1000dm3 = 0.00155 kg/dm3
  • 0.00155kg / 1dm3 x 1000g / 1kg = 1.55g/dm3
  • 1.55g / 1dm3 x 1dm3 / 1L = 1.55 g/L

atomic theory of matter - developed by John Dalton  

• all elements are composed of atoms
• atoms in each element are identical
• chemical reactions can't change, create, or destroy atoms
• compounds are combinations of atoms
• atoms - smallest particles of an element that still contain all the characteristics of that element
• law of constant composition - kinds of atoms and their ratios are constant in compounds
• law of conservation of mass - total mass after a reaction is equal to total mass before the reaction
• law of multiple proportions - atoms combine in compounds in simple ratios

cathode ray tubes - vacuum tube through which high voltage creates radiation  

• produce rays that cause glass to fluoresce, give off light
• electrons move from the cathode (-) to the anode (+)
• J. J. Thomson - summarized his cathode ray observations and found the electron's electrical charge to be 1.76 x 108 coulombs per gram
• Robert Millikan - used his "oil-drop" experiment to calculate the mass of the electron by using his measurements of the charge in a single electron and Thomson's calculations

radioactivity - spontaneous emission of radiation  

• Ernest Rutherford - aka "the second Newton"; discovered the 3 types of radiation
• alpha radiation - positive charge, more massive than beta; equal to nuclei of helium
• beta radiation - negative charge, less massive than alpha; equal to high speed electrons/cathode rays
• gamma radiation - unaffected by electrical fields, has no charge

nuclear atom - final view of atom discovered by Rutherford  

• J. J. Thomson - believed that electrons were embedded in a positively charged sphere
• Ernest Rutherford - used gold foil experiment to prove that positive charge resided in a very small, dense region called the nucleus
• protons discovered by Rutherford in 1919
• neutrons discovered by James Chadwick in 1932

subatomic particles - determine the atom's characteristics  

• electronic charge - 1.602*10-19 coulombs
• atoms have the same number of protons/electrons, no net charge
• atomic mass unit (amu) - used to measure atomic mass; equal to 1.66054 x 10-24 grams, 1/12 the mass of a carbon-12 atom
• angstrom - 10-10 meters; along w/ picometers, used to express atomic diameters;
• atomic diameter equal to about 1-5 angstroms, atomic nuclei equal to about 10^-4 angstroms
• interiors of collapsed stars may approach the density within atomic nuclei

basic forces - 4 basic interactions in nature  

• gravity - attraction between all objects in proportion to their masses; incredibly small between atoms
• electromagnetism - forces between electrically charged objects
• strong nuclear force - acts between subatomic particles in the nucleus to keep them together
• weak nuclear force - weaker than electrical forces, but stronger than gravity; shows up in certain types of radioactivity

isotope - elements that differ in number of neutrons  

• atomic number - number of protons
• atomic mass - number of protons/neutrons
• nuclide - atom of a specific isotope
• average atomic mass - found by using the masses of all its isotopes and their relative abundance; aka atomic weight

periodic table - developed in 1869  

• similar elements placed in the same column (group)
• chemical/physical properties show repetition
• metallic elements - on the left/middle part of the periodic table; shiny, malleable, conductors
• nonmetallic elements - separated from metals on the table by a diagonal line from boron to astatine; brittle, dissolves to form salt, insulators
• metalloids - elements that have properties falling between metals/nonmentals; shiny/brittle
• 8A - unreactive, inert/noble gases; no charge
• 1A - alkali metals; +1 charge
• 2A - alkaline earth metals; +2 charge
• 6A - chalcogens; -2 charge
• 7A - halogens; -1 charge

molecules - combination of 2+ atoms bound tightly together  

• chemical formula - uses chemical symbols and subscripts to show what elements are in the compound
• diatomic molecule - made of just 2 atoms (exists naturally in this state); includes the halogens (exluding astatine), oxygen, hydrogen, nitrogen
• molecular compounds - contains more than 1 type of atom
• empirical formula - shows the subscripts of the molecular formula in the smallest possible ratio
• structural formula - shows how the substances are bonded to each other
• perspective drawing - structural formula written to give a sense of 3D shape
• ball-and-stick model - shows atoms as spheres, bonds as sticks
• space-filling model - shows relative sizes of atoms up in size
• ion - charged particle formed when electrons are added/removed from an atom
• cation - positively charged ion
• anion - negatively charged ion
• polyatomic ions - molecules with a net charge
• charges relate to the atomic position on the periodic table
• ionic compounds - contains both positively/negatively charged ions; generally combinations of metals/nonmetals
• molecular compounds - contains nonmetals only

 
ammonium NH4+ acetate C2H3O2- acetic HC2H3O2 hydronium H3O+ hydrogen carbonate HCO3- carbonic H2CO3   hypochlorite ClO- hypochlorous HClO phosphite PO33- chlorite ClO2- chlorous HClO2 phosphate PO43- chlorate ClO3- chloric HClO3   perchlorate ClO4- perchloric HClO4 carbonate CO32- bromate BrO3- chromic H2CrO4 chromate CrO42- hydroxide OH- hydrobromic HBrO3 dichromate Cr2O72- nitrite NO2- nitrous HNO2 oxalate C2O42- nitrate NO3- nitric HNO3 sulfite SO32- iodate IO3- oxalic H2C2O4 sulfate SO42- permanganate MnO4- phosphorous H3PO3 peroxide O22- cyanide CN- phosphoric H3PO4   thiocyanate SCN- sulfurous H2SO3     sulfuric H2SO4     hydrocyanic HCN

chemical nomenclature - naming of substances  

• some compounds known for a long time have traditional, "common" names (ex. water, ammonia)
• organic compounds - contain carbon
• inorganic compounds - don't contain carbon

naming positive ions  

• cations formed from metal atoms have the same name as the metal
• w/ metals that have different charges, use numerals to show the positive charge
  • Fe2+ iron (II)
  • Fe3+ iron (III)
• "-ous" added to latin form of metal name for ion w/ lower charge
  • Cu+ cuprous
  • Fe2+ ferrous
• "-ic" added to latin form of metal name for ion w/ higher charge
  • Cu2+ cupric
  • Fe3+ ferric
• cations formed from nonmetal atoms have names ending in "-ium"
  • NH4+ ammonium
  • H3O+ hydronium

naming negative ions  

• replace the endings of monoatomic anions w/ "-ide"
  • Cl- chloride
  • O2- oxide
  • N3- nitride
• oxyanions - polyatomic anions containing oxygen; names end w/ "-ate" for the most common form, "-ite" for the version w/ 1 less oxygen
  • SO42- sulfate
  • SO32- sulfite
  • NO3- nitrate
  • NO2- nitrite
• prefixes - "per-" indicates anion w/ additional oxygen than the oxyanion ending in "-ate"; "hypo-" indicates anion w/ 1 less oxygen than the oxyanion ending in "-ite"
  • ClO- hypochlorite
  • ClO2- chlorite
  • ClO3- chlorate
  • ClO4- perchlorate
• add word "hydrogen" or "dihydrogen" to the oxyanion name if H+ was added (thus decreasing the negative charge)
  • CO32- carbonate
  • HCO3- hydrogen carbonate

naming ionic compounds  

• cation name followed by the anion name
  • CaCl2 calcium chloride
• use roman numerals if the cation has multiple possible charges
  • CuClO copper (I) chlorite
  • FeSO4 iron (II) sulfate

naming acids  

• equal to an anion connected to enough hydrogen ions to balance the charge
• add "hydro-" and "-ic" to anion names that end in "-ide"
  • HCl hydrochloric acid
  • H2S hydrosulfuric acid
• use "-ic" for anions ending in "-ate"
  • HClO3 chloric acid
• use "-ous" for anions ending in "-ite"
  • HClO2 chlorous acid

naming binary molecular compounds  

• write the name of the element farthest to the left on the periodic table first (but always write oxygen last)
• write the lower element first if both are in the same group
• 2nd element gets an "-ide" ending
• use Greek prefixes to indicate how much of each element is in the compound (but don't use "mono-" for the 1st element)
  • Cl2O dichlorine monoxide
  • NF3 nitrogen trifluoride

alkanes - most basic class of hydrocarbons, each carbon atom is bonded to 4 other atoms  

• hydrocarbon - compound containing only carbon and hydrogen
• name ends in "-ane"
• long alkanes use greek prefixes to show how many carbon atoms are used
• alcohol - replaces a hydrogen atom w/ an -OH group; name ends in "-ol"
• polyethylene - alkane series extended to include tens of thousands of carbon; used to make plastic products
• polymer - substance made of thousands of smaller molecules

stoichiometry - used to find quantitative information about the substances in reactions  

• based on law of conservation of mass (discovered by Antoine Lavoisier)
• atoms are neither created/destroyed in chemical reactions
• atoms are only rearranged

chemical equations - represents chemical reactions  

• reactants >> products
• reaction must have same number of each atom on both sides to be balanced
  • balanced equation - Cl2 + O >> Cl2O
  • unbalanced equation - H2 + O2 >> H2O
• use only coefficients to balance the equation (don't change subscripts)
• best to always balance the elements that appear in the fewest formulas on each side

combination reaction - 2+ substances react to form 1 product  

• metal/nonmetal combine to form ionic solids
• A + B >> C
  • C + O2 >> CO2
  • CaO + H2O >> Ca(OH)2

decomposition reaction - 1 substance reacts to produce 2+ substances  

• usually involves heating
• C >> A + B
  • CaCO3 >> CaO + CO2
  • 2KClO3 >> 2KCl + 3O2

combustion reaction - rapid reactions that produce a flame  

• usually involves O
• usually forms water and carbon dioxide when hydrocarbons are burned
  • C3H8 + 5O2 >> 3CO2 + 4H2O
  • 2CH3OH + 3O2 >> 2CO2 + 4H2O

formula weight - sum of the atomic weights in each chemical formula  

• equal to atomic weight if finding the weight of an element
• aka molecular weight if finding the weight of a compound
• formula units - represents the chemical formulas of substances; useful to describe ionic compounds (formula unit of ionic compounds same as its empirical formula)
• percent composition - percentage of mass of each element in a substance; equal to number of atoms times atomic weight divided by formula weight
• Find the percent composition of carbon in C2H2 (acetylene)
  • 2(12) / (2 x 12 + 2 x 1) = 24 / 26 = 0.92 = 92%
• Find the percent composition of oxygen in C6H8O6 (ascorbic acid)
  • 6(16) / (6 x 12 + 8 x 1 + 6 x 16) = 96 / 176 = 0.55 = 55%

Avogadro's number (mole) - number of atoms in 12g of pure carbon-12  

• equal to 6.02*10²³
• molar mass - numerically equal to the element's atomic mass (1 atom of carbon-12 = 12 amu, 1 mol of carbon-12 = 12 grams)
• use dimensional analysis to convert between masses, moles, and numbers of particles

moles to number of representative particles  

• multiply number of moles by 6.02 x 10²³
• How many carbon particles in four moles of carbon? 4 x 6.02 x 10²³ = 2.408 x 10²⁴

number of representative particles to moles  

• divide number of representative particles by 6.02 x 10²³
• How many moles in 6.02 x 10²³ atoms? (6.02 x 10²³)/(6.02 x 10²³) = 1

mass (grams) to mole  

• divide the mass by the representative particle's molar mass
• How many moles are in 24 grams of carbon-12? 24/12 = 2

moles to mass (grams)  

• multiply number of moles by the representative particle's molar mass
• How many grams are in 23 moles of iron? 23 x 55.845 = 1.3 x 10³

number of representative particles to mass (grams)  

• divide number of representative particles by 6.02 x 10²³ and then multiply that by the representative particle's molar mass
• What's the mass of 5.234 x 10³⁰ silicon atoms? (5.234 x 10³⁰) / (6.02 x 10²³) x 28.0855 = 2.441 x 10⁸

mass (grams) to number of representative particles  

• divide the mass by the representative particle's molar mass and then multiply it by 6.02 x 10²³
• How many particles of carbon-12 are found in 12 grams of it? 12 g / 12/g/mol x 6.02 x 10²³ = 6.02 x 10²³

empirical formula - tells relative number of atoms of each element in a compound  

• finding the ratios of elements in a compound gives the empirical formula (from percentage composition)
  • C₆H₁₂O₂ (molecular formula) >> CH₂O (empirical formula)
• subscripts in the molecular formula can be found by multiplying the subscripts in the empirical formula by a whole number
• whole number multiple can be found by dividing the molecular weight by empirical formula weight
• combustion analysis - all the C goes into CO₂ and all the H goes into H₂O in combustion; moles of C/H in original compound can be found from the masses of CO₂/H₂O in the product
• Find empirical/molecular formula of caffeine, which contains 49.5% carbon, 5.15% hydrogen, 28.9% nitrogen, 16.5% oxygen. It has a molar mass of 195 g.
  • 49.5 / 12 = 4.125 mol carbon
  • 5.15 / 1 = 5.15 mol hydrogen
  • 28.9 / 14 = 2.1 mol nitrogen
  • 16.5 / 16 = 1.0 mol oxygen
  • C : H : N : O = 4:5:2:1
  • empirical formula is C₄H₅N₂O
  • 49.5% x 195 / 12 = 8.0 mol carbon (2x amount in empirical calculations)
  • molecular formula is C₈H₁₀N₄O₂
• Find empirical/molecular formula for nicotine (contains carbon, hydrogen, nitrogen) if 5.250 mg nicotine combusted to form 14.242 mg CO₂ and 4.083 mg H₂O. It has a molar mass of 160 g.
  • 14.242 + 4.083 - 5.250 = 13.075 mg O₂ used
  • 12 / (12 + 2 x 16) = 27% C in CO₂
  • 27% x 14.242 = 3.85 mg C
  • 2 / (2 + 16) = 11% H in H₂O
  • 11% x 4.083 = 0.45 mg H
  • 5.250 - 0.45 - 3.85 = 0.95 mg N
  • 0.95 / 5.250 = 18% N; 0.45 / 5.250 = 8.6 % H; 3.85 / 5.250 = 74% C
  • 18 / 14 = 1.3 mol N; 8.6 / 1 = 8.6 mol H; 74 / 12 = 6.2 mol C
  • C : H : N = 5 : 7 : 1
  • empirical formula is C₅H₇N
  • 18% x 160 / 14 = 2.1 mol nitrogen (about 2x amount in empirical calculations)
  • molecular formula is C₁₀H₁₄N₂

quantitative information - available from the coefficients of a balanced equation  

• coefficients represent the relative number of molecules/moles of the substance
  • 2H2O same as 2 molecules of water, or 2 moles of water
• substances in a balanced equation are stoichiometrically equivalent
• use dimensional analysis to predict amount of products/reactants
  • In the chemical equation 2H2 + O2 >> H2O, how many moles of water can be created from 4 moles of hydrogen?
  • 4 mol H2 x 1 mol H2O / 2 mol H2 = 2 mol H2O
• combinations of conversion factors can be used to predict amount of products/reactants when given known masses and a balanced equation
  • How much water can be made from 23.1 grams of hydrogen?
  • 23.1 g x 1 mol H2 / 2.01588 g x 1 mol H2O / 2 mol H2 x 18.01528 g = 103 g H2O

limiting reactant - restricts how much of the excess substance can be used  

• determines how much of the product can possibly form
  • When given 23.1 grams of hydrogen and 25.4 grams of oxygen, how much water can be produced?
  • 23.1 g x 1 mol H2 / 2.01588 g x 1 mol H2O / 2 mol H2 x 18.01528 g = 103 g H2O
  • 25.4 g x 1 mol O2 / 31.9988 g x 1 mol H2O / 1 mol O2 x 18.01528 g = 14.3 g H2O
  • oxygen, the limiting reactant, would produce less water product than hydrogen
  • no more than 14.3 grams of water can be made from the reactants given
• theoretical yield - quantity of product predicted to form if all the limiting reactant is used
• percent yield - actual yield divided by theoretical yield
  • What is the percent yield if 12.5 g of water was produced from 23.1 g of hydrogen and 25.4 g of oxygen? (see above)
  • 12.5 / 14.3 = 0.874 = 87.4%

aqueous solution - solutions in which water is the dissolving medium  

• water able to dissolve wide assortment of substances
• most chemical reactions require water to take place
• solution - homogeneous mixture of 2+ substances
• solvent - substance in the solution in greater quantity
• solute - other substances besides the solvent in a solution

electrolytic properties - conducts electricity due to the ions floating in the solution 

• electrolyte - substance whose aqueous solution contains ions
• nonelectrolyte - substance whose aqueous solution doesn't contain ions
• usually only ionic compounds can form electrolytic solutions
• polarity of water breaks apart ionic compounds and keeps them apart
• certain molecular substances (ex. acids) can break into ions
• strong electrolyte - dissociates completely into ions
• weak electrolyte - dissociates partly into ions; ions and original compound exists in a state of chemical equilibrium (reversible reaction)
• double arrow - represents ionization of weak electrolytes
• single arrow - represents ionization of strong electrolytes
• soluble ionic compounds are strong electrolytes

precipitation reaction - reactions that result in the formation of an insoluble product 

• precipitate - insoluble solid formed by a reaction in the solution
• occurs when oppositely charged ions attract each other so strongly that they form an insoluble solid
• solubility - amount of substance that can be dissolved in a given quantity of water
• insoluble - defines any substance w/ solubility less than 0.01 mol/L
• in reactions between 2 strong electrolytes, determine all the possible combinations of ions to see which are soluble

soluble ionic compounds

• all compounds containing nitrate (NO3-)
• all compounds containing C2H3O2-
• all compounds containing ammonium (NH4+)
• all compounds containing alkali metals
• compounds w/ chloride (Cl-) besides thoses w/ Ag+, Hg22+, Pb2+
• compounds w/ bromide (Br-) besides those w/ Ag+, Hg22+, Pb2+
• compounds w/ iodine (I-) besides those w/ Ag+, Hg22+, Pb2+
• compounds w/ sulfate (SO4 2-) besides those w/ Sr2+, Ba2+, Hg22+, Pb2+

insoluble ionic compounds

• compounds w/ sulfide (S2-) besides those w/ Ca2+, Sr2+, Ba2+
• compounds w/ carbonate (CO32-)
• compounds w/ phosphate (PO43-)
• compounds w/ hydroxide (OH-) besides those w/ Ca2+, Sr2+, Ba2+

metathesis (exchange) reactions - reactions where ions exchange partners

• ex. single replacement or double replacement reactions
• occurs in precipitation/acid-base reactions

ionic equations - sometimes useful to show if the dissolved substances are ions or molecules

• molecular equation - shows chemical formulas of reactants/products w/o indicating ionic characteristics
• complete ionic equation - equation written showing all soluble electrolytes as ions
• spectator ions - ions that appear in identical forms on both sides of an equation
• net ionic equation - when spectator ions are left out
• if all ions in a complete ionic equation are spectators, then no reaction occurs
• more than 1 set of reactants can result in the same net ionic equation

acids - substances that ionize to form hydrogen ions

• aka proton donors (H+ is essentially just a proton)
• monoprotic acids - donate 1 hydrogen ion per molecule (ex. HCl, HNO3)
• diprotic acids - donate 2 hydrogen ions per molecule (ex. H2SO4); ionization occurs in 2 steps

bases - substances that accept hydrogen ions, produce hydroxide ions in reactions

• ex. NaOH, KOH, Ca(OH)2
• doesn't necessarily have to contain hydroxide ions
• must accept hydrogen ion

strength of acids/bases

• strong electrolyte = strong acid/base
• weak electrolyte = weak acid/base
• reactivity - depends on the actions of both the cation/anion, not just strength of acid/base
• strong acids - HCl, HBr, HI, HClO3, HNO3, H2SO4
• strong bases - group 1A/2A metal hydroxides
• molecular compounds that aren't acids/bases are nonelectrolytes

neutralization reactions - mixing an acidic and base solution

• product has none of the characteristics of the reactants
• salt - ionic compound whose cation comes from a base and whose anion comes from an acid
• neutralization reaction between an acid and metal hydroxide produces water and a salt
• sulfate/carbonate combine w/ hydrogen ions to form gas

oxidation-reduction (redox) reactions - involve transfer of electrons between reactants

• oxidation - loss of electrons in a substance
• reduction - gain of electrons in a substance
• in reactions, 1 reactant loses an electron and another reaction gains an electron
• oxidation/reduction come together

oxidation number - aka oxidation state; actual chargeof the atom as a monoatomic ion

• oxidation >> increase in oxidation number
• reduction >> decrease in oxidation number
• equals 0 when atom is in its elemental form
• equals charge for all monoatomic ions
• usually negative for all nonmetals
• sum of all oxidation numbers in a neutral compound equals 0

displacement reaction - ion in a solution gets replaced through oxidation of an element

• A + BX >> AX + B
• metals undergo displacement reactions w/ acids to form salts and hydrogen gas
• net ionic reaction shows a change in oxidation states

activity series - list of metals arranged in order of decreasing oxidation

• oxidation causes metals to be eaten away
• active metals - metals at the top of the list; alkali/alkaline earth metals; reacts most readily to form compounds
• noble metals - metals at the bottom of the list; 8B/1B metals
• any metal can be oxidized by elements below it on the list
• only metals above hydrogen in the activity series can react w/ acids to form H2

concentration - amount of solute dissolved in a given quantity of solvent

• greater the solute >> greater the concentration
• molarity = moles of solute / liters of solution
• concentrations of electrolytic solutions depends on chemical formula
  • can be given as molarity of compound or molarity of individual ions
• molarity can be used as conversion factor between volume of solution and moles of solute

dilution - lowers the concentration of a solution by adding water

• stock solutions - solutions in concentrated form
• moles of solute before dilution = moles of solute after dilution
• molarity of concentrated solution * volume of concentrated solution = molarity of diluted solution * volume of diluted solution

chemical analysis - convert to moles to find relationships between reactants/products  

titration - determines concentration of a particular solute in a solution  

• combines sample of solution w/ standard solution (reagent solution of known concentration)
• can be conducted suing acid-base, precipitation, or redox reactions
• equivalence point - point when stoichiometrically equivalent quantities are brought together
• indicators - acid-base dyes that show when the equivalence point is reached
• end point - when the color change of dyes occur; usually very close to equivalence point

energy - ability to do work or transfer heat  

• force - any type of push/pull placed on an object
• work - energy used to make an object move against a force
  • work = force x distance
• heat - energy transferred from hotter object to colder one
• kinetic energy - energy of motion
  • Ek = 1/2 mv2
  • thermal energy - from kinetic energy of molecules in a substance
• potential energy - energy based on relative position
  • Ep = mhg, where g = 9.8 m/s2 (gravitational constant)
  • chemical energy - from potential energy in bonds
• electrostatic energy - interactions w/ charged particles; form of potential energy
  • Eel = (kQ1Q2) / d
  • positive when 2 charges repel each other, negative when attracting each other
• joule - SI unit for energy
• calorie - amount of energy required to raise temperature of 1g of water by 1°
  • 1 calorie = 4.184 joules
  • 1 Cal = 1 nutritional calorie = 1000 calories = 1 kcal

internal energy - sum of all kinetic/potential energy in a system  

• 1st law of thermodynamics - energy is conserved, moves between system/surroundings
• numerical value of E can't be found, only change in E (DE)
• DE = Efinal - Einitial
• DE negative when energy lost (exothermic)
• DE positive when energy gained (endothermic)
• DE = q + w
  • q positive when heat moves from surroundings into system, negative when heat moves from system into surroundings
  • w positive when surroundings work on system, negative when system works on surroundings

system - area isolated for study  

• surroundings - everything other than the system in question
• closed system - can exchange energy but not matter w/ surroundings

state function - property of system determined by its condition/state  

• doesn't depend on how a change is carried out
• examples of state functions - E, DE
• not examples of state functions - work, heat

enthalpy - describes heat flow in chemical changes  

• H = E + PV
• DH = DE + P DV = q + w - w = q = heat change at constant pressure
• DH positive w/ endothermic reactions, negative w/ exothermic reactions
• DH = Hfinal - Hinitial = H(products) - H(reactants)
• enthalpy of reaction - aka heat of reaction, DHrxn; enthalpy change during reaction
• extensive property (depends on moles used in reaction)
• reverse sign of enthalpy if reaction goes in reverse
• depends on state reactants/products (liquid/solid/gas)
  • enthalpy of vaporization (DH for liquid to gas conversion)
  • enthalpy of fusion (DH for solid to liquid conversion)
• enthalpy of combustion (DH for combusting substance in oxygen)

Hess's Law - DH for reaction equals sum of DH for its parts  

• same value of DH regardless of how many steps used for the final reaction

Calculate DH for 2C(s) + H2(g) >> C2H2(g)  

• Given:
  • C2H2 + 5/2 O2 >> 2CO2 + H2O DH = -1299.6 kJ
  • C + O2 >> CO2 DH = -393.5 kJ
  • H2 + 1/2 O2 >> H2O DH = -285.8 kJ
• 2CO2 + H2O >> C2H2 + 5/2 O2 DH = +1299.6 kJ
• 2C + 2O2 >> 2CO2 DH = 2(-393.5) = -787.0 kJ
• 2C + 2O2 + H2 + 1/2 O2 + 2CO2 + H2O >> C2H2 + 5/2 O2 + 2CO2 + H2O DH = 1299.6 - 787.0 - 285.8 = 226.8 kJ
• 2C + 2O2 + H2 + 1/2 O2 + 2CO2 + H2O >> C2H2 + 5/2 O2 + 2CO2 + H2O DH = 1299.6 - 787.0 - 285.8 = 226.8 kJ
• 2C + 2O2 >> C2H2 DH = 226.8 kJ

enthalpy of formation - aka heat of formation, DHf  

• standard state - pure form at atmospheric pressure, 25°C
• standard enthalpy - enthalpy change when reactants/products in standard state
• stardard enthalpy of formation - change in enthalpy for reaction that makes 1 mol of a substance from its elements (everything in standard state)
  • DH°rxn = SDH°f(products) - SDH°f(reactants)
  • equal to 0 for substances in elemental form

Find standard enthalpy change for combustion of C6H6(l) into CO2(g) and H2O(l)  

• Given:
  • DH°f(CO2(g)) = -393.5 kJ/mol
  • DH°f(H2O(l)) = -285.8 kJ/mol
  • DH°f(C6H6(l)) = 49.0 kJ/mol
• C6H6 + 15/2 O2 >> 6CO2 + 3H2O
• DH°rxn = SDH°f(products) - SDH°f(reactants)
• = [6mol x DH°f(CO2) + 3mol x DH°f(H2O)] - [1mol x DH°f(C6H6) + 15/2 mol x DH°f(CO2)
• = [6mol (-393.5 kJ/mol) + 3mol (-285.8 kJ/mol)] - [1mol (49.0 kJ/mol) + 15/2 mol (0 kJ/mol)]
• = -2361 - 857.4 - 49.0 = -3267 kJ

calorimetry - measurement of heat flow  

• calorimeter - device used to measure heat flow
• heat capacity - amount of heat needed to raise temperature by 1°C
  • molar heat capacity - heat capacity of 1 mol of a substance
• specific heat - heat capacity of 1g of a substance
  • specific heat = quantity of heat transferred / (grams of substance x temperature change)
  • = q / (m x DT)
• q = specific heat x grams of substance x DT

Find amount of heat needed to warm 250 g of water from 22°C to 98°C  

• Given:
  • specific heat of water = 4.18 J/g-K
• change in K same as change in °C
• q = specific heat of water x grams of water x DT
• = 4.18 J/g-K x 250g x 76K
• = 7.9 x 104 J

constant-pressure calorimetry - heat gained by solution same as heat lost by reaction  

• qsoln = specific heat of solution x grams of solution x DT = -qrxn

Find enthalpy change for reaction if temperature of solution in calorimeter changed from 21.0°C to 27.5°C when 50mL of 1.0M HCl was mixed with 50mL of 1.0M NaOH  

• Given:
  • total volume of solution = 100 mL
  • density of solution = 1.0g/mL
  • specific heat of solution = 4.18 J/g-K
• mass of solution = 100mL x 1.0g/mL = 100g
• temperature change = 27.5 - 21.0 = 6.5°C = 6.5K
• qrxn = -(4.18 J/g-K)(100g)(6.5K) = -2.7 x 103 J = -2.7 kJ
• 50mL x 1M = (0.050L)(1.0 mol/L) = 0.050 mol HCl or NaOH in solution
• enthalpy change per mol = -2.7 kJ / 0.050 mol = -54 kJ/mol

bomb calorimetry (constant-volume calorimetry) - studies combustion reactions  

• compound reacts w/ excess oxygen
• combustion started by electrical spark
• qrxn = -Ccal x DT
  • Ccal = heat capacity of calorimeter

Find heat of reaction for comubstion of 4.00g of CH6N2 if its combustion in a calorimeter w/ heat capacity of 7.794 kJ/ °C changes temperature from 25.00°C to 39.50°C  

• qrxn = -Ccal x DT
• DT = 39.50 - 25.00 = 14.50°C
• Ccal = 7.794 kJ/°C
• qrxn = -(7.794 kJ/°C)(14.50°C) = -113.0 kJ
• heat of reaction per mol = -113.0 kJ / 4.00g CH6N2 x 46.1g CH6N2 / 1 mol CH6N2 = -1.30 x 103 kJ/mol CH6N2

fuel value - energy released when 1g of material is combusted  

• proteins/carbohydrates produce about 17kJ/g
• fats produce 38kJ/g
• fossil fuels - formed from decomposition of plants/animals
  • natural gas - gaseous hydrocarbons
  • petroleum - liquid containing mostly hydrocarbons
  • coal - solid w/ hydrocarbons of high molecular weight; most abundant fossil fuel
  • syngas - "synthesis gas"; coal undergoes coal gasification to become mixture of methane, hydrogen, carbon dioxide gas
• nuclear energy - energy released through splitting/fusion of atom nuclei
• renewable energy - essentially inexaustible energy
  • solar energy from sun w/ solar cells (photovoltaic devices)
  • wind energy from windmills
  • geothermal energy from heat in Earth's mass
  • hydroelectric energy from flowing rivers
  • biomass energy from crops, biological waste

electromagnetic radiation - aka radiant energy  

• visible light is a type of electronmagnetic radiation
• speed of light = 3.00 x 108 m/s
• wavelength - distance between successive peaks/troughs
• frequency - # of complete wavelengths/cycles that pass a point per second
• amplitude - intensity of radiation; max height of wave
• (frequency)(wavelength) = speed of light = c = nl
  • frequency = n
  • wavelength = l
• smaller wavelength >> higher frequency >> more powerful radiation

Planck's constant - 6.63 x 10-34 J-s  

• quantum - "fixed amount;" smallest quantity of energy that can be emitted/absorbed
• energy always emitted in chunks of minimum size
• energy = Planck's constant x frequency
  • E = hn
• energy quantized, can only increase in steps (though steps are extremely small)
• photoelectric effect - photons striking metals cause electrons to be emitted
  • photon - energy packet, transfers energy to electrons
  • energy of photon = E = hn
• 6.02 x 1023 photons per mol

Find energy of 1 photon light w/ wavelength 589 nm  

• c = nl
• 3.00 x 108 m/s = n(589 nm)
• n = (3.00 x 108 m/s) / (5.89 x 10-7 m) = 5.09 x 1014 s-1
• E = hn
• E = (6.63 x 10-34 J-s)(5.09 x 1014 s-1) = 3.37 x 10-19 J/photon

line spectra - spectrum containing radiation of specific wavelengths  

• monochromatic radiation - consists of a single wavelength
• spectrum - separation of radiation into different wavelengths
• continuous spectrum - contains light of all wavelengths

Rydberg equation - allowed calculation of wavelengths of all spectral lines  

• 1/l = (Rh)(1/n12 - 1/n22) = [(2.18 x 10-18 J) / hc] (1/n12 - 1/n22)
• Rh = 1.096776 x 107 m-1

Bohr's Model - electrons moving in circular paths lose energy and spiral towards nucleus  

• only orbits w/ certain radii, dependent on energies of electrons
• electron in allowed energy state has specific energy, doesn't radiate energy
• energy emitted/absorbed by electrons when it changes energy states
• E = (-2.18 x 10-18 J)(1/n2) = energy in hydrogen atom
  • n = integer from 1 to ¥ = quantum number
  • ground state - lowest energy state
  • excited state - higher energy state
  • E = (-2.18 x 10-18 J)(1/¥2) = 0
• DE = Efinal - Einitial = Ephoton = hn = hc/l = (-2.18 x 10-18 J)(1/nf2 - 1/ni2)
  • ni = initial energy state
  • nf = final energy state
  • l = hc / DE
  • n = = DE / h
• doesn't explain spectra of any atom besides hydrogen
• electrons actually show properties of waves

Find the de Broglie wavelength of an electron w/ velocity 5.97 x 106 m/s  

• Given:
  • l = h/(mv)
  • m = 9.11 x 10-28 g = 9.11 x 10-31 kg
  • h = 6.63 x 10-34
  • v = 5.97 x 106
• l = (6.63 x 10-34) / (9.11 x 10-31 x 5.97 x 106)
• l = 1.22 x 10-10 m

matter waves - describes wave characteristics of material particles  

• all matter have wavelengths
• wavelength inversely proportional to mass/velocity
• l = Planck's constant / (mass x velocity) = h / mv
• microscopes use wave properties of electrons to map out surfaces

quantum (wave) mechanics - studies w/ subatomic particles  

• wave functions - mathematical functions describing electron's matter wave
  • represented by y, but has no physical meaning by itself
  • probability density = y2
  • electron density - areas w/ high probability of finding electrons
• orbital - describes distribution of electron density in space
  • not the same as the orbits mentioned in Bohr's model
  • path of electron can't be precisely tracked (due to uncertainty principle)
• quantum numbers - n, l, ml used to describe orbitals
  • n = positive integer, principle quantum number; increases as electron gets farther away from nucleus
  • En = (-2.18 x 10-18 J)(1/n2)
  • l = integer from 0 to n-1, azimuthal quantum number; describes orbital shape
  • ml = integers between l and -l, magnetic quantum number; describes orientation of orbital
• electron shell - collection of orbitals w/ same n value
• subshell - orbitals w/ same n and l values
• n subshells in the shell w/ principle quantum number n
• each subshell has 2(l)+1 number of orbitals
• n2 = total number of orbitals in a shell
• ground state - electron in lowest energy orbital
• excited state - electron in any other orbital, reaches this level by absorbing photons

orbitals  

• s orbital - lowest energy orbital, has spherical symmetry
  • y² approaches 0 as it gets farther apart from nucleus
  • nodes - regions where y² goes to 0
  • orbital size increases as n increases
• p orbital - has dumbbell shaped orbital w/ 2 lobes
  • shape shows average distribution of electrons, not its path
  • every shell beyond the 1st has 3 p orbitals, each w/ different spatial orientation
• d orbital - occur beyond 1st 2 shells
  • 5 different d orbitals
• f orbital - occur beyond 1st 3 shells
  • 7 different f orbitals
• in atoms w/ multiple electrons, repulsions cause different subshells to be at different energies
• degenerate orbitals - orbitals w/ same energy

Pauli exlusion principle - no 2 electrons can have same set of quantum numbers (n, l, m_l, m_s)  

• electron spin - property of electrons where each electron spins on an axis
• spin generates magnetic field
• m_s = +1/2 or -1/2, spin magnetic quantum number
• each orbital can hold 2 electrons, as long as they spin in opposite directions

Hund's rule - in degenerate orbitals, lowest energy obtained if max # of electrons have same spin  

• electrons w/ same spin magnetic quantum number have parallel spins
• electrons naturally repel one another, tend to occupy different orbitals if they can

electron configuration - way electrons are distributed among atoms' orbitals  

• ground state - electrons in lowest possibly energy states
• orbitals filled in order of increasing energy (max 2 electrons per orbital)
• electrons w/ opposite spins are paired when in same orbital, unpaired electrons do not have an electron w/ opposite spin in same orbital
• condensed electron configuration - abbreviated electron configuration
  • aka noble gas notation
  • electron configuration of nearest noble gas w/ lower atomic number represented by chemical symbol in brackets
  • only outer electrons matter in reactions
  • core electrons - inner-shell electrons
  • valence electrons - electrons given after noble gas core
• lanthanide (rare earth) elements - fill up the 4f orbitals
• actinide elements - fill up the 5f orbitals
• representative (main-group) elements - in s and p blocks of periodic table

• potassium (K) >> 1s²2s²2p⁶3s²3p⁶4s¹ or [Ar]4s¹
• iron (Fe) >> 1s²2s²2p⁶3s²3p⁶4s²3d⁶ or [Ar]4s²3d⁶

periodic table development - ordered according to atomic mass by Mendeleev/Meyer  

• Mendeleev got majority of credit for advancing his ideas more vigorously and predicting the existence of substances yet to be discovered
• Henry Moseley - arranged atoms by atomic number instead of atomic mass, solved the problems in the original periodic table

effective nuclear charge - electric field created by nucleus and surrounding electron density  

• uses average environment created by nucleus/electrons
• Zeff = Z - S
  • Z = number of protons in the nucleus
  • S = average number of core electrons
• inner electrons shield outer electrons from the nucleus' charge
• charge increases as you move across any row/period of the periodic table
  • Z increases as S stays the same, so Zeff increases
• charge increases only slightly as you move down a column/family
  • larger electron cores less able to shield outer electrons than smaller cores

atomic radii - nonbonding/bonding radius  

• nonbonding radii (van der Waals radii) - radius of atoms not in molecules
• bonding radii (covalent radii) - radius of atoms in bonds
  • slightly shorter than nonbonding radii
• increases as you move down column/family
  • outer electrons get farther from nucleus as principal quantum number increases
• decreases as you move left to right across a row/period
  • as Zeff increases, outer electrons get pulled in more
• ionic radii - depends on charge
  • cations - smaller than neutral atoms
  • anions - larger than neutral atoms
• isoelectronic series - series of ions w/ same number of electrons, different number of protons
  • radius decreases as nuclear charge increases

ionization energy - measures amount of energy needed to lose an electron  

• more difficult to remove electrons w/ greater ionization energy
• more energy needed to remove each subsequent electron
• sharp increase in ionization energy needed to remove inner shell electrons
  • inner shell electrons much closer to nucleus
• I1 generally increases w/ atomic number on each row
• I1 decreases as atomic number increases down a group
• representative elements have larger range of I1 than transition elements
• smaller atoms tend to have higher ionization energies (electrons closer to nucleus)

ion electron configurations - electrons removed from largest available quantum number first  

• electrons added to lowest available quantum number first

electron affinity - energy change when electron is added to a gaseous atom  

• measures attraction of atom for added electron
• usually negative (energy usually released when electron is added), but can be positive for noble gases (anion higher in energy than separated atom/electron)
• halogens have the most negative electron affinities
• noble gases have positive electron affinities (when adding an electron would place it on a new energy subshell)
• group 5A (w/ 1/2 filled subshells) have electron affinities either positive or less negative than group 4A
• doesn't change much going up/down a group

metals - make up 3/4 of the periodic table, in left/middle part  

• shiny luster
• malleable (can be pounded into thin sheets)/ductile (can be drawn into wires)
• good conductors of heat/electricity
• tend to form cations (have low ionization energies) in aqueous solutions
• compounds of metals w/ nonmetals tend to be ionic
• most metal oxides are basic (form hydroxides when dissolved in water)
• metal oxides w/ acid form salt
• alkali metals (group 1A) - soft metallic solids
  • have the lowest I1 on each row
  • hydride ion - H-; bonds w/ alkali metals to form hydrides
  • reacts exothermically w/ water
  • superoxide - O2-; combines w/ potassium, rubidium, cesium
  • emit certain colors when heated by a flame
• alkaline earth metals (group 2A) - denser/harder than alkali metals
  • less reactive than alkali metals
  • Mg used for lightweight structure alloys because layer of MgO protects it from other chemicals

nonmetals - on right side of periodic table  

• no luster, has various appearances
• poor conductors of heat/electricity
• most nonmetal oxides are acids
• tend to form anions/oxyanions in aqueous solutions
• molecular substances - compounds made up only of nonmetals
• hydrogen - doesn't really belong in a particular group
  • exists as H2 gas in most conditions
  • can be metallic at very high pressures
  • has much higher I1 than other alkali metals (lacks any type of nuclear shielding)
• oxygen group (group 6A)
  • elements change from nonmetal to metal as you go down the group
  • allotrope - different form of same element in same state
  • ozone - O3, less stable than O2
  • sulfur - exists naturally as S8 rings
• halogens (group 7A) - aka "salt formers"
  • melting/boiling points increase as atomic number increases
  • very highly negative electron affinities
  • Cl - most industrially useful halogen
  • forms halide compounds w/ hydrogen
• noble gases (group 8A) - monoatomic nonmetals, gas at room temperature
  • have completely filled s and p subshells
  • very unreactive

metalloids - have properties intermediate between metals/nonmetals  

• have only some metallic properties, but lack others
• many used as electrical semiconductors, integrated circuits

chemical bond - attraction between atoms or ions

• ionic bond - electrostatic forces existing between ions of opposite charge
  • ions formed when electrons transferred
  • ionic compounds made from metals and nonmetals
• covalent bond - sharing of electrons
  • interactions between nonmetallic elements
• metallic bond - attraction between metals
  • each atom bonded to many neighboring atoms
  • bonding electrons free to move throughout the substance

ionic bonding - ions held in 3D array

• in forming ionic compounds, 1 atom loses an electron while another gains 1
• formation of ionic compounds usually exothermic
  • ionic compounds stable due to attraction of opposite charges
  • energy released by ion attraction makes up for endothermic ionization
• lattice energy - energy required to completely separate a mole of ionic compound into its ions
  • increases as charges on ions increase, radii decrease
• makes ionic compounds very hard/brittle, w/ high melting points
• covalent bonds hold atoms in polyatomic ions together, but polyatomic ions as a whole still act as ions
• E = k(Q1Q2)/d
  • Q1, Q2 = charges on the particles
  • d = distance between centers
  • k = constant 8.99 x 109 J-m/C2
  • energy increases as charge increases, or as distance between centers decrease

electron configurations of ions of representative elements

• by octet rule, ions tend to have electron configurations of noble gases
• electrons lost from subshell w/ highest n value first

electron configurations of metal ions

• usually only have net charges of +1, +2, or +3
• tries to have a full d subshell, loses electrons in s subshell first (higher n value)

covalent bonding - sharing of electron pairs

• nuclei’s attraction to the shared electrons overcomes their repulsion to each other
• single line used in Lewis structures to represent each shared electron pair
  • single bond - only 1 pair of electrons shared
  • double bond - 2 pairs of electrons shared, represented by 2 lines
  • triple bond - 3 pairs of electrons shared
  • multiple bonds shorter/stronger than single bonds
• # of shared electron pairs increase >> distance between bonded atoms decrease
• metals w/ high oxidation numbers tend to act molecular instead of ionic

electronegativity - ability of atom in a molecule to attract electrons

• used to see if bond will be nonpolar covalent, polar covalent, or ionic
• bond polarity - describes sharing of electrons between atoms
  • nonpolar covalent bond - electrons shared equally (electronegativities equal)
  • polar covalent bond - 1 atom attracts bonding electrons more than the other
  • ionic bond forms when difference in electronegativity > 3
• based on ionization energy, electron affinity
• fluorine - most electronegative
• cesium - least electronegative
• in molecules, electron density (and negative charge) shift towards more electronegative atom

dipole moment - measure of dipole’s magnitude

• dipole - formed when electrical charges of opposite sign, equal magnitude separated by distance
• m = Qr
  • m = dipole moment
  • Q = electronic charge (unit e = 1.60 x 10-19 C)
  • r = distance separating 2 atoms
• debyes (D) - unit equal to 3.34 x 10-30 coulomb-meters (C-m)

Calculate the dipole moment of HCl if the H and Cl are separated by 1.27Å 

• Q = 1e = 1.60 x 10-19 C
• r = 1.27Å
• m = (1.60 x 10-19 C)(1.27Å)(10-10m/1.00Å)(1D/3.34 x 10-30 C-m)
• m = 6.08 D

In units of e, what is the charge on the atoms of a compound if they're separated by 2.74Å and have a measured dipole moment of 1.97 D? 

• Q = m/r
• m = 1.97D = (1.97D)(3.34 x 10-30 C-m / D) = 6.5798 x 10-30 C-m
• r = 2.74Å = (2.74Å)(10-10m/1.00Å) = 2.74 x 10-10m
• Q = m/r = (6.5798 x 10-30 C-m) / (2.74 x 10-10m) = 2.40 x 10-20 C
• (2.40 x 10-20 C)(1 e / 1.60 x 10-19 C) = 0.150 e

drawing Lewis structures -

• valence electrons - electrons involved in chemical bonding
• dots placed on 4 sides of chemical symbol
• number of valence electrons in representative elements same as group number
• add all valence electrons
  • same # as group # for representative elements
  • add electron for each negative charge in anions
  • subtract electron for each positive charge in cations
  • only keep track of total number of electrons (no need to associate electrons to certain atoms)
• write symbols for atoms
  • central atom usually the least electronegative
  • single bond (2 shared electrons) represented by dash between atomic symbols
• complete octets of atoms
  • arrange dots so all atoms have 8 electrons
  • hydrogen only needs 2 in outer shell
• place extra electrons around central atom
  • even if it adds up to more than 8
• use multiple bonds if not enough electrons
  • use unshared electron pairs to make double/triple bonds

formal charge - charge of an atom if all atoms had the same electronegativity

• tells how electrons are distributed in a molecule
• # of valence electrons - # of electrons assigned to the atom
  • # of assigned electrons = all nonbonding electrons + 1/2 of bonding electrons
• sum of formal charges = overall charge of molecule

resonance structures - placement of electrons changes, but atom placement doesn’t

• double arrow used to show resonance structures
• cannot be described by a single Lewis structure
• actual structure an average of the possible Lewis resonance structures

exceptions to octet rule -

• odd # of electrons - not possible for octet to form around all atoms
  • not possible for complete pairing of electrons
• less than an octet - weird exception for boron/beryllium
  • usually due to halogen’s reluctance to form double bonds
• more than an octet
  • extra electrons placed around central atom
  • expanded electron shells only possible for elements in period 3 and beyond
  • d subshell can be used for bonding
  • central atom usually bonded to smallest/most electronegative atoms

Draw Lewis Structures for PCl3, PO33-, XeFl4, BCl3 

• PCl3
• PO33-
• XeFl4
  • more than an octet
• BCl3
  • less than an octet

bond enthalpy - enthalpy change for breaking of a bond

• strength of covalent bond determined by amount of energy needed to break it
• always positive value
• DHrxn = (bond enthalpies of bonds broken) - (bond enthalpies of bonds formed)
• derived for gaseous molecules, only averaged values

Find the DH for 2C2H6 + 7O2>> 4CO2 + 6H2O  

• C2H6
• CO2
• H2O
• Given bond enthalpies:
  • C-H 413 kJ/mol
  • C-C 348 kJ/mol
  • C=O 799 kJ/mol
  • H-O 463 kJ/mol
  • O2 495 kJ/mol
• 2 C-C, 12 C-H, 7O2 bonds broken
• 8 C=O, 12 H-O bonds created
• [2(348) + 12(413) + 7(495)] - [8(799) + 12(463)] = -2831 kJ

Find the DH for N2H4 >> N2 + 2H2  

• N2H4
• N2
• Given bond enthalpies:
  • N-H 391 kJ/mol
  • H-H 436 kJ/mol
  • N=N 941 kJ/mol
  • N-N 163
• 1 N-N, 4 N-H bonds broken
• 1 N=N, 2 H-H bonds created
• [163 + 4(391)] - [941 + 2(436)] = -86 kJ

VSEPR model - valence-shell electron-pair repulsion model

• bond angles - determines shape of molecules
• electron domain - area where electrons are most commonly found
  • bonding pair - electrons between 2 atoms in bonds
  • nonbonding pair - lone pair of electrons not shared
• nonbonding pairs, multiple bonds have greater repulsive forces on nearby electron domains than single bonds
• best arrangement minimizes repulsions between electron domains
• electron-domain geometry - determines arrangement of electron domains
• molecular geometry - determines arrangement of atoms
• atoms beyond 3rd period can have over 4 pairs of electrons around it (has extra d subshell)

bond polarity - measure of electron sharing

• overall dipole = sum of bond dipoles
  • dipoles have magnitude/direction
  • consider shape to find overall polarity
• nonpolar molecule - overall dipole moment of 0
  • bond dipoles cancel each other out
• polar molecule - nonzero overall dipole moment
  • when molecule not the same all around

valence-bond theory - explains covalent bonding

• orbitals overlap to form covalent bonds
• optimum distance between bonded nuclei
• bond length - distance where attractive forces (between electrons/nuclei) are balanced by repulsive forces (between electrons/electrons, nuclei/nuclei)

hybrid orbitals - new orbitals formed by a mix of other orbitals

• hybridization - process of mixing/changing atomic orbitals
• energy released by bond formation > energy needed to promote electrons to higher levels
• sp hybrid orbital - 1 electron from s subshell promoted to p subshell
  • able to form 2 bonds
  • forms linear arrangement
• sp2 hybrid orbital - 1 electron from s subshell promoted to p subshell
  • 1 electron in s subshell, 2 electrons in p subshell
  • forms trigonal planar arrangement
• sp3 hybrid orbital - 1 electron from s subhsell promoted to fill p subshell
  • forms tetrahedral arrangement
• hybridization w/ d orbitals - possible w/ atoms in 3 rd period and beyond
  • sp3d hybrid orbital - trigonal bipyramidal arrangement
  • sp3d2hybrid orbital - octahedron arrangement

multiple bonds - overlap of p orbitals

• sigma bonds - single bonds where electron density concentrated symmetrically around line connecting the nuclei
• pi bonds - overlapping electron regions lie above/below internuclear axis
  • can’t experimentally be seen
• single bond - 1 sigma bond
• double bond - 1 sigma bond, 1 pi bond
• triple bond - 1 sigma bond, 2 pi bonds
• delocalized pi bonding - cannot be described as individual electron bond between atoms
  • split up all around the molecule
  • explains identical length of alternating single/double bonds in molecules like benzene
  • found in all molecules w/ 2+ resonance structures

molecular orbital theory - describes electrons in terms of molecular orbitals 

• bonding molecular orbital - lower energy molecular orbital
  • electron attracted to both nuclei
  • very stable
• antibonding molecular orbital - little electron density between nuclei
  • higher energy molecular orbital
  • atomic orbitals cancel each other out, area of largest electron density found on opposite sides of nuclei
• energy-level diagram - aka molecular orbital diagram
  • shows interactions of atomic orbitals
• bond order = 1/2(# of bonding electrons - # of nonbonding electrons)
  • 0 >> no bond
  • 1 >> single bond
  • 2 >> double bond
  • 3 >> triple bond
• # of molecular orbitals = # of atomic orbitals combined
• atomic orbitals combine best w/ same type of atomic orbitals
• overlap increase >> bonding lower in energy, antibonding higher in energy
• each molecular orbital can hold at most 2 electrons (w/ alternate spins)
• antibonding/bonding pi molecular orbitals lower in energy than antibonding sigma orbitals, higher in energy than bonding sigma orbitals (unless interactions between s and p levels exist)

molecular properties of electron configurations

• paramagnetism - molecules w/ unpaired electrons attracted into magnetic fields
• diamagnetism - molecules w/ no unpaired electrons weakly repelled by magnetic fields
• bond order increase >> bond distance decrease, bond enthalpy increase
• heteronuclear diatomic molecule - diatomic molecule made up of different atoms
  • homonuclear diatomic molecule - diatomic molecules made up of same atoms
  • will resemble homonuclear molecular orbital diagram if electronegativities are not too far off
  • may have non-integer bond order value

characteristics of gas -  

• possible for substance to coexist as solid, liquid, and gas at the same time
• vapor - gaseous form of a substance normally existing as liquid/solid
• expands to fill the container it’s in (gas volume = volume of container)
• pressure added to gas >> gas gets compressed easily >> volume decreases
• form homogeneous mixtures (regardless of type of gases involved)

pressure - force acted upon a given area

• P = F/A
• newton (N) - SI unit for force, kg-m/s2
• pascal (Pa) - SI unit for pressure, N/m2
  • kPA = 1000 Pa
  • bar = 105 Pa, approximate atmospheric pressure at sea level
• barometer - invented by Evangelista Torricelli, measures atmospheric pressure
  • manometer - used to measure pressure of liquids/gases, similar to barometer
• standard atmospheric pressure = pressure needed to support column of mercury 760 mm high
  • atm = 760 mm Hg, 760 torr = 1.01325 x 105 Pa = 101.325 kPa
  • 1 torr = 1 mm Hg

Boyle’s Law - pressure-volume relationship

• volume of gas (at constant temperature) inversely proportional to pressure
• P1V1 = P2V2

Charle’s Law - temperature-volume relationship

• volume of gas (at constant pressure) directly proportional to temperature
• V1/T1 = V2/T2

Avogadro’s Law - quantity-volume relationship

• law of combining volumes - at given pressure/temperature, volumes of reacting gases exist in simple ratios
• equal volumes of gases at same temperature/pressure have equal numbers of molecules
• volume of gas (at constant pressure/temperature) directly proportional to number of moles of gas
• V = constant x n

ideal-gas equation - describes hypothetical gas (ideal gas)

• results not exactly correct for actual gases
• PV = nRT
  • R = gas constant, depends on values of P, V, n, T
  • T = temperature, always expressed as absolute temperature
  • n = number of moles of gas
  • P = pressure, usually given in atm
  • V = volume, usually given in L
• standard temperature and pressure (STP) - 0 C, 1 atm
  • molar volume - 22.41 L/mol
• all gas laws derived from the ideal-gas equation
• P1V1/T1 = P2V2/T2

Find the temperature of gas at which 0.407 mol takes up 3.23 L of space at 118 in Hg 

• Given:
  • PV = nRT
  • P = 118 in = (118 in x 2.54 cm/in x 10 mm/cm) = 2997.2 mm = 3.94 atm
  • V = 3.32 L
  • n = 0.407
  • R = 0.08206 (L-atm/mol-K)
• (3.94)(3.23) = (0.407)(0.08206)T
• T = [(3.94)(3.23)] / [(0.407)(0.08206)]
• T = 381 K

gas density - unit mass per unit volume

• n/V = P/(RT)
• density = n(molar mass)/V = P(molar mass)/(RT)
• higher molar mass/pressure >> higher gas density
• higher temperature >> lower gas density

Find the density of carbon tetrachloride at 714 torr at 125° C  

• Given:
  • molar mass = 12 + 4(35.5) = 154 g/mol
  • P = 714 torr = 714/760 atm
  • R = 0.0821 L-atom/mol-K)
  • T = 125° C = 125 + 273 K = 398 K
• d = [(714/760)(154)] / [(0.0821)(398)]
• d = 4.43 g/L

Dalton ’s law of partial pressures

• total pressure of a mixture of gases equal to sum of pressures of each gas in the mixture
• Pt = P1 + P2 + P3 +…
• at constant temperature/volume, total pressure determined by number of moles of gas
• mole fraction - ratio of moles of 1 substance in mixture to total number of moles
• partial pressure = mole fraction x total pressure

Find the partial pressures and total pressure of a mixture made from 6.00 g O2 and 9.00 g CH4 in a 15.0 L container at 0° C  

• Given:
  • PV = nRT
  • P = nRT/V
  • R = 0.0821 (L-atm/mol-K)
  • T = 273 K
• moles of O2 = 6.00 / 32 = 0.188 mol
• moles of CH4 = 9.00 / 16 = 0.563 mol
• pressure of O2 = (0.188)(0.0821)(273) / 15 = 0.281 atm
• pressure of CH4 = (0.563)(0.0821)(273) / 15 = 0.841 atm
• total pressure = 0.281 + 0.841 = 1.122 atm

collecting gases over water - gases produced by chemical reactions often collected over water

• volume of gas measured by raising/lowering inverted bottle until water level is the same on inside/outside of bottle
• pressure inside bottle equal to atmospheric pressure outside when water level is the same
• Ptotal = Pgas + PH2O

kinetic-molecular theory - explains behavior of gases

• gases made up of large number of molecules in constant, random motion
• gas molecules are extremely tiny
• attractive/repulsive forces between gas molecules don’t really do anything
• energy transferred between molecules during collisions, but average kinetic energy stays the same (elastic collisions)
  • individual molecules have different speeds
• average kinetic energy directly proportional to absolute temperature
  • constant temperature >> constant average kinetic energy
• rms speed (u) - speed of molecule w/ average kinetic energy
  • average kinetic energy = 1/2 mu2
  • u = (3RT/molar mass)1/2

effusion - escape of gas molecules through a tiny opening

• Graham’s law - compares rates of effusion under identical conditions
  • shows that lighter gas effuses more rapidly
  • r1/r2 = (M2/M1)1/2
• diffusion - spread of substance throughout a space or 2nd substance
  • molecular collisions >> motion of gas molecules constantly changes >> diffusion is slow
  • mean free path - average distance a molecule moves between collisions

behavior of real gases - different than behavior of ideal gases

• ideal gas equation assumes that gas molecules take up no space, have no intermolecular forces
• less deviation w/ higher temperature, lower pressure
• gas volumes usually slightly greater than predicted by ideal-gas equation
• gas pressure usually slightly lesser than predicted by ideal-gas equation

van der Waals equation - takes into account the gas volume and attractive forces

• (P + n2a/V2)(V - nb) = nRT
• constants a, b different for each gas

Find the pressure in atm that O2 exerts at 70.6° C if 1.850 moles occupies 16.5 L.  

• Given:
  • a = 1.36 (L2 atom/ mol2)
  • b = 0.0318 (L/mol)
  • R = 0.08206 (L-atm/mol-K)
  • (P + n2a/V2)(V - nb) = nRT
• (P + (1.850)2(1.36)/(16.5)2)(16.5 - (1.850)(0.0318)) = (1.850)(0.08206)(343.6)
• P = [(1.850)(0.08206)(343.6)] / [(16.5 - (1.850)(0.0318))] - (1.850)2(1.36)/(16.5)2
• P = 3.16 atm

3 phases - dependent on the intermolecular forces

• gas
  • takes on volume/shape of container
  • compressible
  • flows easily
  • diffusion occurs rapidly
• liquid
  • takes on shape of container
  • doesn’t expand to fill container
  • incompressible
  • flows readily
  • diffusion occurs slowly
• solid
  • keeps its own shape/volume
  • incompressible
  • doesn’t flow
  • diffusion occurs very slowly

intermolecular forces - much weaker than ionic/covalent bonds

• molecules remain intact when intermolecular forces broken
• stronger intermolecular forces >> higher melting/boiling points
• van der Waals forces - intermolecular attractive forces between neutral molecules
  • dipole-dipole, London dispersion, hydrogen bonding
• ion-dipole force - between ion and partial charge on an end of polar molecular
  • ion charge increase or dipole moment magnitude increase >> increase in magnitude of attraction
  • important for solutions of ionic substances in polar liquids
• dipole-dipole force - attraction between positive/negative ends of neutral polar molecules
  • only effective when polar molecules very close together
  • increasing polarity >> increasing intermolecular attractions
  • smaller molecules >> increasing attraction
• London dispersion force - instantaneous dipole moment
  • found between nonpolar and polar molecules
  • due to temporary dipole moments at particular moments
• polarizability - how easily an electric field can change the molecule’s charge distribution
  • increasing molecular weight >> increasing strength of dispersion forces
  • most massive/polar molecules have largest attraction forces
• hydrogen bonding force - intermolecular attraction between hydrogen atom in polar bond and unshared electron pair on another molecule
  • generally stronger than dipole-dipole/dispersion forces
  • hydrogen’s small size lets it get close to electronegative atom
  • equal amounts of water take up more volume as solid than as liquid (usually reversed for other substances)

phase changes - changes of the state of a substance

• vaporization - liquid to gas
• condensation - gas to liquid
• melting - solid to liquid
• freezing - liquid to solid
• deposition - gas to solid
• sublimation - solid to gas
• loss of movement >> loss of energy >> exothermic
• more movement >> gain of energy >> endothermic
• heat of fusion - energy needed to change solid to liquid
  • H_fus
• heat of vaporization - energy needed to change liquid to gas
  • usually larger than heat of fusion since molecules must get rid of all intermolecular attractions when in transition from liquid to gas
  • H_vap
• heating curve - temperature of system versus amount of heat added
  • temperature stops at the points where substance changes states
  • supercooling - temporarily cooling a liquid below its freezing point w/o it turning into a solid
• critical temperature - highest temperature at which a liquid phase can form
• distinct pressure - pressure needed to turn gas into liquid at the critical temperature

Find the enthalpy change as 10 g of a liquid at 70° C goes to a gas at 100° C  

• Given:
  • liquid boils at 90° C
  • specific heat of liquid = 1.0 J/g-K
  • specific heat of gas = 0.3 J/g-K
  • enthalpy of vaporization = 8.5 J/g
• enthalpy of heating liquid from 70° C to 90° C = (10)(1.0)(20) = 200 J
• H_vap = (10)(8.5) = 85 J
• enthalpy of heating gas from 90° C to 100° C = (10)(0.3)(10) = 30 J
• total enthalpy change = 200 + 85 + 30 = 415 J

vapor pressure - pressure exerted by vapor during dynamic equilibrium

• weaker attractive forces >> more molecules able to escape into gas phase >> higher vapor pressure
• dynamic equilibrium - when evaporation/condensation occur at same rate
  • no net change in system
  • doesn’t occur in an open container
• volatility - describes liquids that evaporate readily
  • higher vapor pressure >> more easily evaporates
• normal boiling point - boiling point of liquid at 1 atm
  • occurs when vapor pressure equals external pressure on liquid’s surface
  • higher pressure >> higher boiling point

phase diagrams - graphic representation of equilibrium between phases

• uses temperature/pressure as axis
• lines represent equilibrium between phases
• triple point - where 3 curves intersect
  • all 3 phases are in equilibrium at this point

liquids - properties explained by intermolecular forces

• viscosity - resistance of a liquid to flow
  • measured by timing to see how long it takes an amount of liquid to flow through a tube
  • poise (P) - unit of viscosity, g/cm-s
  • lower temperature, larger molecular weight >> lower viscosity
• surface tension - energy needed to increase surface area by a certain amount
  • cohesion - forces that bind similar molecules to each other
  • adhesion - forces that bind substance to a surface
• capillary action - rise of liquids up narrow tubes
  • used by plants to move water/nutrients upwards
  • adhesion between liquid and tube increases surface area

solid structures - either crystalline or noncrystalline (amorphous)

• crystalline solid - molecules arranged in well-defined arrangements
  • have flat surfaces, definite angles, regular shapes
  • brings particles in closest contact >> maximizes attractive forces
  • unit cells - repeating unit of a crystalline solid
  • crystal lattice - 3D array of points representing the crystal
  • primitive cubic - unit cell where lattice points at corners only
  • centered cubic - additional lattice point at center of cell
  • face-centered cubic - additional lattice points at center of each face
• packing of spheres - each sphere surrounded by 6 others in each layer
  • spheres rest in depressions of surrounding layers
  • hexagonal close packing - 3rd layer repeats the 1st, 4th layer repeats the 2nd
  • cubic close packing - 4th layer repeats the 1st
  • coordination number - number of particles immediately surrounding 1 particle on all sides
• amorphous solid - no orderly structure, usually made of large complicated substances

bonding in solids - arrangement of particles determines melting point, hardness, etc

• molecular solids - have molecules/atoms held together by intermolecular forces
  • soft, low melting points
• covalent-network solids - atoms held together in large networks by covalent bonds
  • much stronger than molecular solids, w/ higher melting points
• ionic solids - ions held together by ionic bonds
  • strength depends on charges of ions
  • charges/relative sizes determine structure of solid
• metallic solids - made entirely of metal atoms
  • each atom has 8 or 12 surrounding atoms
  • valence electrons delocalized throughout entire solid
  • mobility of electrons promotes conductivity

solution process - like molecules dissolve like molecules

• ionic compounds dissolved in polar solvent (w/ ion-dipole forces)
• covalent compounds dissolved in nonpolar solvent (w/ dispersion forces)
• forces between solvent/solute must be greater than forces between solute molecules
• solvation - dissolving solute w/ solvent (solvent molecules completely surround solute)
  • hydration - solvation when solvent is water
• Hsoln = H1 + H2 + H3
  • H1 = separation of solute molecules (endothermic)
  • H2 = separation of solvent molecules (endothermic)
  • H3 = forming solvent-solute interactions (exothermic)
  • Hsoln can be either exothermic/endothermic
• no solution if too endothermic, spontaneous reaction if exothermic

spontaneous solution formation - usually exothermic

• energy decrease >> reaction starts spontaneously
• entropy/disorder increase >> reaction starts spontaneously (even if endothermic)
• molecules unrestrained >> spontaneous mixing occurs
• no solution of solute-solute or solvent-solvent forces greater than solvent-solute forces
• solute can change or remain unchanged after solvation

crystallization - opposite of solvation

• saturated solution - will not dissolve more solute if added
• solubility - amount of solute needed to form saturated solution
• unsaturated solution - dissolves less solute than in saturated solution
• supersaturated solution - contains more solute than needed for saturation
  • possible at different temperatures
  • very unstable, will crystallize w/ just a little bit of added solute

solute-solvent interactions - determines tendency of substances to mix

• stronger solute-solvent interaction >> greater solubility
• polar liquids dissolve polar solutes, don’t dissolve nonpolar solutes
• miscible - describes substances that dissolve in each other
• immiscible - describes substances that don’t dissolve in each other
• increase # of polar groups >> increase solubility in water

pressure effects - doesn’t affect liquid/solid as much as gas

• increase pressure over solvent >> increase solubility of gas
• gas solubility increases directly proportional to partial pressure above the solution
• Henry’s Law - Sg = kPg
  • Sg = solubility of gas in solution phase
  • Pg = partial pressure of gas over solution
  • k - Henry’s law constant

temperature effects - different for gas/solid

• temperature increase >> solid solubility increase, gas solubility decrease
• thermal pollution - higher temperatures >> lower oxygen solubility in lakes

concentration - dilute (small solute concentration), concentrated (large solute concentration)

• mass percentage = mass of component in solution / total solution mass
  • parts per million = mass percentage x 106
  • parts per billion = mass percentage x 109
• mole fraction - moles of component in solution / total moles of all components
  • molarity (volume) = moles of solute / liters of solution
  • molality (mass) = moles of solute / kilograms of solvent

Find the mass percentage of a KCl solution if it's supposed to be isotonic w/ a 9% NaCl solution 

• assume the mass of solution = 100 g
• 9% NaCl solution >> 9 g of NaCl
• 9g NaCl x 1 mol NaCl / 58.5g NaCl = 0.1538 mol NaCl >> 0.3076 mol solute
• to be isotonic, KCl solution should have same amount of solute
• 0.3076 mol solute >> 0.1538 mol KCl
• 0.1538 mol KCl x 74.5g KCl / 1 mol KCl = 11.5g KCl
• 11.5 / 100 = 11.5% KCl by mass

In a certain ore, there is 1g of Silver for every 500kg of ore. What is the concentration in parts per million? parts per billion?  

• 500kg = 500,000g
• 1 / 500,000 x 106 = 2 parts per million
• 1 / 500,000 x 109 = 2000 parts per billion

Find the molality of chloride ions if 1 g of CaCl2 is dissolved in 750 mL of water  

• for every molecule of CaCl2, 2 Cl- ions are produced when dissolved
• 1g CaCl2 x 1 mol CaCl2 / 111g CaCl2 = 0.009 mol CaCl2 >> 0.018 mol Cl-
• m = mol solute / kg solvent = 0.018 / 0.75 = 0.024

Find the molarity of H2SO4 in a 95% by mass solution if the solution's density is 1.84 g/cm3  

• assume mass of solution = 100 g
• 95% solution >> 95g H2SO4
• 95g H2SO4 x 1 mol H2SO4 / 98g H2SO4 = 0.97 mol H2SO4
• 1.84 g/cm3 = 1.84 g/mL
• 100g solution / 1.84 g/mL = 54.35 mL solution
• M = mol solute / liters solution = 0.97 / 0.05435 = 17.8

colloids (colloidal dispersions) - in between heterogeneous mixtures and solutions

• solutes larger than molecules, not so large that they separate due to gravity
• large enough to scatter light (Tyndall effect)
• hydrophilic (water loving) and hydrophobic (water fearing) colloids in human body
  • hydrophilic groups shield hydrophobic groups from water
  • hydrophobic colloids prepared in water if ions adsorbed (adhered) onto surface
• emulsifying agent - has hydrophilic and hydrophobic end, helps w/ digestion by shielding hydrophobic substances from water
• coagulation - enlarges colloids so they can be removed from solutions
  • caused by heating or adding electrolyte
• semipermeable membranes also remove colloids (ions can pass through, but colloids can’t)

colligative properties - depends on quantity of solute, not the type

• Raoult’s law - adding solute to solvent lowers vapor pressure
  • partial pressure of solvent vapor = mole fraction of solvent x vapor pressure of pure solvent
  • limited to nonvolatile/nonelectrolyte substances (ideal solution)
• boiling point increases when solute added to solution
  • more solute >> decreased vapor pressure >> takes longer to reach atmospheric pressure (boil)
  • DTb = Kbm
• freezing point decreases when solute added to solution
  • DTf = Kfm
• osmosis - net mov’t of solvent moves towards area w/ higher solute concentration
  • liquid continually moves across semipermeable membrane in both directions
  • osmotic pressure - pressure needed to prevent osmosis
  • isotonic - solutions w/ same osmotic pressure
  • hypotonic - solution w/ lower osmotic pressure
  • hypertonic - solution w/ higher osmotic pressure
  • crenation - shriveling of (hypotonic) cell when liquid moves out
  • hemolysis - rupturing of (hypertonic) cell when liquid moves in
• more solute added >> boiling point increase, freezing point decrease, vapor pressure decrease, salinity increase

Arrange the following solutions from lowest to highest freezing point, lowest to highest boiling point, and lowest to highest vapor pressure: 0.35 m antifreeze, 0.20 m KBr, 0.20 m K2CO3, 0.12 m NaCl, 0.20 m sugar  

• for colligative properties, only number of particles matter
• certain molecules split into ions when dissolved
• 0.35 m antifreeze, 0.20 m sugar don't split into ions
• 0.20 m KBr >> splits into K+, Br- >> 0.40 m solute solute
• 0.20 m K2CO3 >> splits into 2K+, CO32- >> 0.60 m solute
• 0.12 m NaCl >> splits into Na+, Cl- >> 0.24 m solute
• vapor pressure - more solute >> lower vapor pressure
  • 0.20 m K2CO3, 0.20 m KBr, 0.35 m antifreeze, 0.12 m NaCl, 0.20 m sugar
• boiling point - more solute >> higher boiling point
  • 0.20 m sugar, 0.12 m NaCl, 0.35 m antifreeze, 0.20 m KBr, 0.20 m K2CO3
• freezing point - more solute >> lower freezing point
  • 0.20 m K2CO3, 0.20 m KBr, 0.35 m antifreeze, 0.12 m NaCl, 0.20 m sugar

How many grams of NaCl should be added to 3kg of water to get a water-salt solution that freezes at -10 C?  

• Given:
  • DTf = Kfm
  • Kf = 1.86
• water normally freezes at 0 C >> DTf = 10
• 10 = 1.86m
• m = 10 / 1.86 = 5.38
• 5.38 = mol solute / kg water = mol solute / 3
• mol solute = 3 x 5.38 = 16.14 mol solute
• each NaCl provides 2 solutes >> 16.14 mol solute = 8.07 mol NaCl
• 8.07 mol NaCl x 58.5g NaCl / 1 mol NaCl = 472g NaCl

What is the boiling point of a solution w/ 44.4 g CaCl2 in 2L water?  

• Given:
  • DTb = Kbm
  • Kb = 0.51
• CaCl2 splits into 3 solutes in water
• 44.4g CaCl2 x 1 mol CaCl2 / 111g CaCl2 = 0.4 mol CaCl2 >> 1.2 mol solute
• m = 1.2 mol solute / 2kg water = 0.6
• DTf = (0.51)(0.6) = 0.31
• water normally boils at 100 C >> CaCl2 solution boils at 100+0.31 = 100.31 C

Find the molar mass of tabitol if 2.5g of tabitol in a 100mL solution produces an osmotic pressure of 1.79atm at 25C.  

• Given:
  • osmotic pressure = MRT
  • R = 0.0821
  • temperature in kelvin
• 1.79 = mol tabitol / 0.1 L x 0.0821 x (273+25)
• mol tabitol = 1.79 / 0.0821 / 298 x 0.1 = 0.0073
• mol tabitol = 2.5g / molar mass
• molar mass = 2.5 / 0.0073 = 342 g/mol

factors affecting reaction rates -

• chemical kinetics - study of reaction speed
• physical state of reactants - more collisions >> faster reaction
  • solid surface area increase >> reaction rate increase
• reactant concentration - higher concentration >> faster reaction
• temperature - higher temperature >> higher kinetic molecular energy >> more collisions >> faster reaction
• catalyst - substances that increase reaction rate w/o being used up
• collisions must include enough energy and correct positioning to lead to reaction
• average rate = change in concentration / change in time
• instantaneous rate - rate at a specific moment in reaction
  • rates tend to decrease as reaction continues
  • initial rate - reaction rate when reaction first begins
• rate of reactant disappearance = rate of product appearance

rate law - shows how rate depends on concentrations

• aA + bB >> cC + dD
• rate = k [A]m[B]n
• k = rate constant, changes w/ temperature (units of rate / units of concentration2)
• exponents m, n = reaction order
• usually 0, 1, or 2
• shows how concentration affects rate (0 = no change when concentration changed)
• overall reaction order - sum of orders for each reactant
• determined experimentally, not from equation/coefficients

first-order reaction - reaction where rate depends on concentration of single reactant

• ln[A]t = -kt + ln[A]0
• y = mx + b
• only needs 3 quantities to solve for the 4th

second-order reaction - reaction where rate depends on concentrations of 2 reactants  

• 1 / [A]t = kt + 1/[A]0

For the following data and the reaction X + Y >> Z, what is the order for X and Y?  
initial X concentration initial Y concentration initial rate 0.200 0.200 7.50 0.400 0.200 30.00 0.200 0.800 30.00

• as X increases by a factor of 2, the rate increases by a factor of 4
  • 4 = 22
  • for X, reaction is second order
• as Y increases by a factor of 4, so does the rate
  • rate and Y concentration increases at same rate
  • for Y, reaction is first order

For a second order reaction, the rate constantis 25 L/mol-s at 20 C. Find the time it takes for the concentration to go from 0.025 M to 0.010 M  

• Given:
  • 1 / [A]t = kt + 1/[A]0
  • k = 25
  • [A]t = 0.010
  • [A]0 = 0.025
• 1 / 0.01 = 25t + 1 / 0.025
• 25t = 100 - 40 = 60
• t = 60/25 = 2.4 sec

For a first order reaction involving popcorn, 6 kernels pop every 5 seconds when there are 150 kernels. How long until 75 of the kernels pop?  

• Given:
  • ln[A]t = -kt + ln[A]0
  • k = rate / concentration
  • [A]t = 75
  • [A]0 = 150
• rate = 6/5
• concentration = 150
• k = 6/5 / 150 = 1 / 125
• ln (75) = -1/125 t + ln (150)
• ln (75) - ln (150) = -1 / 125 t
• t = 87 sec

half-life (t1/2) - time needed for concentration of reactant to drop to 1/2 or original value 

• fast reaction >> short half-life
• t1/2 = -ln (1/2) / k = 0.693 / k for 1st-order reactions (no dependence on initial concentration)
• t1/2 = 1 / k[A]0 for 2nd-order reactions (dependence on initial concentration)

Find the half-life of a substance that decomposes by 20% after 5 years. 

• 0.8 = (1)(1/2)5/x
• ln(0.8) = 5/x ln(1/2)
• ln(0.8) / ln(1/2) = 5/x
• x = 5 ln(1/2) / ln(0.8)
• 15.5 years

Find the age of a piece of wood whose carbon-14 count is 35/min, when a new piece of wood has a count of 125/min.  

• Given:
  • half life of carbon-14 = 5715 years
  • ln[A]t = -kt + ln[A]0
  • [A]t = 35
  • [A]0 = 125
  • t1/2 = -ln (1/2) / k
• 5715 = -ln(1/2) / k
• k = -ln(1/2) / 5715 = 0.00012
• ln(35) = -(0.00012)t + ln(125)
• ln(35) - ln(125) = -(0.00012)t
• t = (ln35 - ln125) / -0.00012
• 10608 years

Find the half life of a substance if 95% of it disappears after 10 years.  

• 0.05 = (1/2)10/x
• ln (0.05) = 10/x (ln(1/2))
• ln (0.05) / ln (1/2) = 10 / x
• x = 10 ln(1/2) / ln(0.05)
• 2.3 years

collision model - based on kinetic-molecular theory

• shows effects of both temperature/concentration on molecular level
• assumes that molecules must collide to react w/ each other
• not all collisions lead to reactions
• orientation factor - molecules need to be in a certain position to react when colliding

activation energy (Ea) - minimum amount of energy needed for reaction to occur

• kinetic energy of colliding molecules used to break bonds
• activated complex (transition state) - atomic arrangement at the point of highest energy
• lower activation energy >> faster reaction rate
• f = e-Ea/RT
  • f = fraction of molecules w/ energy equal to or greater than Ea
  • R = gas constant
  • T = absolute temperature
• Arrhenius equation - takes into account 3 factors (activation energy, # of collisions, fraction of collisions w/ correct orientation)
  • k = Ae-Ea/RT
  • A = frequency factor constant
• ln (k1/k2) = Ea/R (1/T2 - 1/T1)
  • must knowk2 from T2 (another temperature) to calculate k1 and T1)

reaction mechanisms - process by which reaction occurs

• can describe the order in which bonds are broken/formed
• elementary steps (elementary processes) - single event/step
  • rate law based on molecularity for each elementary step
• molecularity - describes # of reactant molecules in an elementary step
  • unimolecular - single molecule involved (rearrangement)
  • bimolecular - collision of 2 molecules
  • termolecular - simultaneous collision of 3 molecules (very rare)
• multi-step mechanism - sequence of elementary steps
  • elementary steps add up to give overall chemical process
  • intermediate - substance that’s not reactant/product in overall reaction
  • rate-determining (limiting) step - slowest step, limits overall reaction rate
  • rate law of overall reaction = rate law of slowest step

catalyst - substance that changes reaction rate

• doesn’t change during reaction
• lowers overall activation energy
• homogeneous catalyst - in same phase as reactants
• heterogeneous catalyst - in different phase than reactants
  • often consists of metals or metal oxides
  • adsorption - binding of molecules to surface (1st step)
  • active sites - where reactants adsorb
• enzymes - biological catalysts
  • usually large protein molecules
  • lock-and-key model - substrates fit into enzymes at specific location w/ certain shape
  • turnover number - # of catalyzed reactions occurring at an active site

equilibrium in reactions - when forward/reverse reactions occur at same rate  

• ratio of partial pressures, concentrations equals constant
• pressures no longer change
• reaction continues to occur, but ratio stays the same (no net change)
• can be reached from either direction

Haber Process - synthesizing ammonia from hydrogen/nitrogen 

• world’s main source for fixed nitrogen
• N2 + 3H2 <<>> 2NH3

equilibrium constant - only changed by temperature, not by reaction mechanism 

• for reaction aA + bB <<>> cC + dD
• Keq = (C concentration)c(D concentration)d / (A concentration)a(B concentration)b
• value determined by coefficients in chemical equation
• value greater than 1 >> equilibrium towards product side
• value less than 1 >> equilibrium towards reactant side
• for any reaction, constant equal to reciprocal of constant for reverse reaction
• constant raised to power equal to number by which reaction is multiplied
• constants multiplied together for net reactions involving 2+ steps
• do not include pure solids, pure liquids, or solvents in finding constant

For the reaction CaCo3(s) >> CaO(s) + CO2(g), K = 0.0108. Find the equilibrium weight of CO2 if 2g of CaCo3, 2g of CaO, and 0.5g of CO2 were placed in a 1 liter flask.  

• don't take account of solids in the reaction
• K = [CO2] = 0.0108 mol
• 0.0108 mol = grams CO2 / CO2 molar mass
• 0.0108 = grams CO2 / (12 + 16 x 2)
• grams CO2 = 0.0108 (44)
• 0.475 g

At equilibrium for the reaction 2NO(g) + Cl2(g) >> 2NOCl, K = 51. If [NOCl] = 0.10 and [Cl2] = 0.20, find [NO]  

• K = 51 = [NOCl]2 / ([NO]2[Cl2])
• [NO]2 = 0.12 / 0.2 / 51 = 0.00098
• [NO] = 0.031

Le Châtelier’s Principle - equilibrium position changed when outside force disturbs the system

• affected by concentration, temperature, pressure
• change in reactant/product concentration >> reaction shifts to use up the added substance
• lower volume, higher pressure >> reaction shifts to produce less moles of gas
• higher volume, lower pressure >> reaction shifts to produce more moles of gas
• higher temperature >> reaction shifts to use up extra energy (depends on whether reaction is exothermic/endothermic)
• catalysts - only changes how fast equilibrium is achieved, not characteristics of equilibrium itself

Describe the effect on CaCO3(s) >> CaO(s) + CO2(g) when:  

• pressure increased
  • shifts so that less gas is produced
  • shifts left
• volume increased
  • same as decreasing the pressure
  • shifts right to produce more gas
• adding CaCO3
  • no effect, it's a solid

For reaction PCl5(g) >> PCl3(g) + Cl2(g), K = 0.015 and DH = -375. What happens when Cl2 is removed?  

• disregard K or DH
  • they have nothing to do w/ the problem
• reaction shifts right to produce more Cl2

For endothermic equation N2(g) + O2(g) >> 2NO(g), K = 4.0 x 10-4. What happens when pressure is increased?  

• again, disregard K
• same amount of moles of gas on both sides of equation
• reaction doesn't shift to either side

For endothermic equation N2(g) + O2(g) >> 2NO(g), K = 4.0 x 10-4. What happens when temperatre is increased?  

• N2(g) + O2(g) + energy >> 2NO(g)
• adding temperature >> adding energy
• reaction shifts right to use up energy

definition - increases the concentration of H+ ions when dissolved in water 

• the H+ ion generally forms H3O+ w/ water
• Bronsted-Lowry definition - donates proton to another substance
  • even applies to substances not dissolved in water
  • weak H-X bond, stable base, positively polarized H >> stronger acid
• Lewis definition - electron-pair acceptor
• pH = -log [H+]
• strong acid - complete dissociates, [H+] = [acid]
  • HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4
• weak acid - partially dissociates
  • Ka = [H+][conjugate base] / [acid]
  • [H+] = [conjugate base] when Ka is very small
• polyprotic acids - have more than 1 ionizable H atom
  • easier to remove 1st proton than second
  • Ka becomes smaller
• binary acids - contains hydrogen and 1 other element
  • bond strength determines acid strength
  • element increases in size >> bond strength decreases >> acid strength increases
  • element electronegativity increases >> polarity increases >> acid strength increases
• oxyacids - OH and oxygen bonded to a central atom
  • element electronegativity increases, # of oxygen increases >> acid strength increases
• carboxylic acids - contains a carboxyl group, largest cateogry of organic acids
  • # of electronegative atoms increase >> acid strength increases

Find the pH of a 0.001 M solution of acetic acid  

• Ka = 1.8 x 10-5
• Ka = [H+] [acetate] / [acetic acid] = x2 / 0.001
• x = [(1.8 x 10-5) (0.001)]1/2 = 1.3 x 10-4 = [H+]
• pH = -log [H+] = 3.87

Find the pH of 0.1 M solution of HCl  

• strong acid, completely dissociates
• pH = -log [HCl] = 1

definition - increases the concentration of OH- ions when dissolved in water  

• Bronsted-Lowry definition - accepts proton from another substance
  • even applies to substances not dissolved in water
• Lewis definition - electron-pair donor
• pOH = -log [OH-]
  • pOH = 14 - pH
• strong base - completely dissociates, [OH-] = [base]
  • hydroxides of alkali metals and heavy alkaline earth metals
• weak base - partially dissociates
  • Kb = [OH-][conjugate acid] / [base]
  • [OH-] = [conjugate acid] when Kb is very small

Find the pH of 0.0030 M solution of ammonia  

• Kb = 1.8 x 10-5
• Kb = [OH-] [ammonium] / [ammonia] = x2 / 0.0030
• x = [(1.8 x 10-5) (0.0030)]1/2 = 2.3 x 10-4 = [OH-]
• pOH = -log [OH-] = 3.63
• pH = 14 - 3.63 = 10.37

conjugate acid-base pairs - HX + H2O >> X- + H3O+  

• HX is acid
  • conjugate acid of X- (add proton)
• H2O acts as base
  • conjugate base of H3O+ (subtract proton)
• X- is base
  • conjugate base of HX (subtract proton)
• H3O+ acts as acid
  • conjugate acid of H2O (add proton)
• conjugate acid of strong base = weak acid
• conjugate acid of weak base = strong acid
• conjugate base of strong acid = weak base
• conjugate base of weak acid = strong base
• Ka x Kb = Kw
  • Kw = 1.0 x 10-14 (water autoionizes)
  • Ka = Kb >> neutral solution

Find the pH for 0.25 M of sodium lactate  

• conjugate acid is lactic acid
• Ka of lactic acid = 1.4 x 10-4
• Kb = Kw / Ka = 10-14 / (1.4 x 10-4) = 7.14 x 10-11
• Kb = [OH-] [lactic acid] / [lactate]
• [OH-] = [(7.14 x 10-11) (0.25)]1/2 = 4.2 x 10-6
• pOH = -log [OH-] = 5.37
• pH = 14 - 5.37 = 8.63

hydrolysis - ability of ions to react w/ water to make H+, OH- ions  

• anions act as bases (adding proton >> acid)
  • identify conjugate acid, use Kw to find the strength
  • anions of strong acids have negligible effect
  • anions of weak acids increase pH
• cations act as acids (subtract proton >> base)
  • identify conjugate base, use Kw to find the strength
  • cations of strong bases have negligible effect
  • cations of weak bases decrease pH

common-ion effect - ionization of electrolyte decreases if common ion added  

• shifts equilibrium against a certain side
• Ka = [H+][common ion] / [acid]
• Kb = [OH-][common ion] / [base]
• can decrease the solubility of certain slightly soluble salts

Find the pH in a solution of 0.1 M lactic acid and 0.1 M sodium lactate  

• Given:
  • Ka of lactic acid = 1.4 x 10-4
  • Ka = [H+][common ion] / [acid]
• 1.4 x 10-4 = [H+] x 0.1 / 0.1
• 1.4 x 10-4 = [H+]
• pH = -log [H+] = 3.85

Find the pH of a 500mL solution containing 1.00 g propionic acid and 1.0 g sodium propionate  

• Given:
  • Ka of propionic acid = 1.3 x 10-5
  • propionic acid = HC3H5O2
  • sodium propionate = NaC3H5O2
  • Ka = [H+][common ion] / [acid]
• [propionic acid] = 1 / (1+36+5+32) / 0.5 = 0.027
• [sodium propionate] = 1 / (23+36+5+32) / 0.5 = 0.0208
• 1.3 x 10-5 = [H+] [sodium propionate] / [propionic acid]
• [H+] = 1.68 x 10-5
• pH = -log[H+] = 4.77

complex ions - made up of metal ion and Lewis base  

• stability depends on equilibrium constant for its formation
• metal ions act as electron-pair acceptors, Lewis acids
• metal salts can dissolve more easily when metal ion can react w/ the solvent

amphoterism - able to act as acid or base  

• includes hydroxides, oxides of Al3+, Cr3+, Zn2+, Sn2+
• contains basic anions >> dissolves in acidic solutions
• can form complex ions >> dissolves in basic solutions

precipitation/separation of ions - occurs when product of ion concentrations > Ksp  

• equilibrium = saturation
• selective precipitation - uses reagent to form a precipitate w/ 1 of the ions
• insoluble chlorides - formed from Ag+, Hg22+, Pb2+
• acid-insoluble sulfides - CuS, Bi2S3, CdS, PbS, HgS, As2S3, Sb2S3, SnS2
• base-insoluble sulfides/hydroxides - formed from Al3+, Cr3+, Fe3+, Zn2+, Ni2+, Co2+, Mn2+
• insoluble phosphates - formed from Mg2+, Ca2+, Sr2+, Ba2+
• alkali meta ions - remainder of possible metal ions, can be tested for individually

buffered solutions - contains weak conjugate acid-base pair  

• resists drastic pH changes
• able to neutralize both H+ and OH- ions
• conjugate acid-base pair can't consume each other
• usually created by combining a weak acid/base w/ its salt
• pH of buffer = log ([base] / [acid]) - log (Ka)
  • pH = pKa + log ([X-] / [HX])
• buffer capacity - amount of acid/base the buffer can neutralize before pH begins changing
• add strong acid >> [HX] increases, [X-] decreases
• add strong base >> [X-] incresaes, [HX] decreases

If the buffer system in blood is made up of carbonic acid and sodium bicarbonate, what is the base to acid ratio if the pH is measured to be 7.41?  

• Given:
  • pH = 7.41
  • Ka of carbonic acid = 4.3 x 10-7
  • Ka = [H+] [X-] / [HX]
• [H+] = 10-7.41
• 4.3 x 10-7 = [H+] [HCO3-] / [H2CO3]
• [HCO3-] / [H2CO3] = 4.3 x 10-7 / 10-7.41 = 11

How many moles of NaOBr should be added to 1.00 L of 0.050 M HOBr to form a buffer w/ pH 8.80?  

• Given:
  • pH = 8.80
  • Ka HOBr = 2.5 x 10-9
  • Ka = [H+] [X-] / [HX]
• assume that the NaOBr has negligible effect on the volume
• [H+] = 10-8.8
• 2.5 x 10-9 = [H+] [NaOBr] / [HOBr]
• [NaOBr] = (2.5 x 10-9) (0.050) / (10-8.8) = 0.079

In a buffer of acetic acid and sodium acetate where the pH is 5, find the molarity of sodium acetate if the molarity of acetic acid is 0.10  

• Given:
  • pH = 5
  • [acetic acid] = 0.10
  • Ka = 1.8 x 10-5
  • Ka = [H+] [X-] / [HX]
• [H+] = 10-5
• 1.8 x 10-5 = [H+] [sodium acetate] / [acetic acid]
• [sodium acetate] = (1.8 x 10-5) (0.1) / (10-5) = 0.18

titration - known concentration of base is added to acid (or acid to base)  

• equivalence point - where amount of acid/base are stoichiometrically equivalent
• pH titration curve - graph of pH as a function of volume of added titrant
• strong acid-strong base titration - characterized by sudden, very steep change in pH
  • initial pH based on initial concentration of acid
  • pH increases slowly when base first added
  • pH increases dramatically near equivalence point
  • only the salt exists at the equivalence point, pH = 7
  • pH after equivalence point determined by amount of base in solution
• weak acid-strong base titration - less change at equivalence point than strong acid-strong base
  • higher initial pH than strong acid-strong base
  • equivalence point always > 7 due to strong base

Find the pH of solution formed when 45.0 mL of 0.100 M NaOH is added to 50.0 mL of 0.100 M acetic acid.  

• Given:
  • Ka = 1.8 x 10-5
• mol of NaOH = 0.045 L (0.100 mol / 1 L) = 4.5 x 10-3 mol
• mol of acetic acid = 0.050 L (0.100 mol / 1 L) = 5 x 10-3 mol
• HC2H3O2 + NaOH >> NaC2H3O2 + H2O
• 5 x 10-3 - 4.5 x 10-3 = 5 x 10-4 mol of acetic acid not consumed by NaOH after reaction
• 4.5 x 10-3 mol of acetate after reaction
• 45.0 mL + 50.0 mL = 0.095 L solution
• [acetic acid] = 5 x 10-4 / 0.095 = 0.0053
• [acetate] = 4.5 x 10-3 / 0.095 = 0.0474
• Ka = [H+] [acetate] / [acetic acid]
• [H+] = (1.8 x 10-5) ( 0.0053) / (0.0474) = 2.0 x 10-6
• pH = -log [H+] = 5.70

Find the pH at the equivalence point in the titration of 50.0 mL of 0.100 M acetic acid w/ 0.100 M NaOH  

• Given:
  • same molarity >> same volume needed to neutralize >> 50.0 mL of NaOH used at equivalence point
  • no acetic acid will be left over at the equivalence point
  • Ka = 1.8 x 10-5
• mol of acetic acid at beginning = 0.0500 L (0.100 mol / 1 L) = 5 x 10-3 = mol of acetate at point
• [acetate] = 5 x 10-3 mol / 0.100 L = 0.05 M
• Kb = Kw / Ka = 10-14 / 1.8 x 10-5 = 5.6 x 10-10
• Kb = [OH-] [acetic acid] / [acetate]
• [OH-] = [(5.6 x 10-10) (0.05)]1/2 = 5.3 x 10-6
• pOH = -log [OH-] = 5.28
• pH = 14 - 5.28 = 8.72

When titrating a 0.1 M acetic acid solution with 0.005 M NaOH, what is the pH 1/2 way through the titration?  

• Given:
  • 1/2 of acetic acid will have reacted w/ NaOH
  • Ka = 1.8 x 10-5
• for comparison's sake, assume there's 100 mL of acetic acid
• mol of acetic acid = 0.100 L (0.1 mol / 1 L) = 0.01 mol
• mol of NaOH used = 0.005 (1/2 of acetic acid)
• volume of NaOH used = 0.005 mol / 0.005 M = 1 L
• HC2H3O2 + NaOH >> NaC2H3O2 + H2O
• 0.005 mol acetic acid left after reaction
• 0.005 mol acetate left after reaction
• [acetic acid] = 0.005 mol / 1.1 L = 0.0045
• [acetate] = 0.005 mol / 1.1 L = 0.0045
• Ka = [H+] [acetate] / [acetic acid]
• [H+] = 1.8 x 10-5
• pH = -log [H+] = 4.74

solubility product - Ksp  

• equals product of ion concentrations raised to the power of their coefficients
• can be used to calculate solubility (g/L)

Find the Ksp for a saturated solution of Mg(OH)2 if the pH is 10.17  

• assume that Mg(OH)2, a strong base, dissociates completely
• Mg(OH)2 >> Mg2+ + 2OH-
• pOH = 14 - pH = 14 - 10.17 = 3.83
• [OH-] = 10-pOH = 10-3.83 = 1.48 x 10-4
• [Mg2+] = [OH-] / 2 = 7.4 x 10-5
• Ksp = [Mg2+] [OH-]2 = 1.62 x 10-12

Find the molar solubility of Mn(OH)2 if the Ksp is 1.6 x 10-13  

• Mn(OH)2 >> Mn2+ + 2OH-
• Ksp = [Mn2+] [OH-]2
• let c = molar solubility, concentration of Mn
• Ksp = c x (2c)2 = 4c3
• 1.6 x 10-13 = 4c3
• c = 3.4 x 10-5

Find the molar solubility of Ba(IO3)2 in a solution of 0.010 M NaIO3  

• Given:
  • Ksp = 6 x 10-10
• [IO3-] = [NaIO3] = 0.010
• Ksp = [Ba2+] [IO3-]2
• 6 x 10-10 = [Ba2+] (0.010)2
• [Ba2+] = molar solubility = 6 x 10-6

Find the molar solubility of Ba(IO3)2 in a solution of 0.01 M NaNO3  

• Given:
  • Ksp = 6 x 10-10
• NaNO3 has no effect on the solubility of Ba(IO3)2
• Ksp = [Ba2+] [IO3-]2
• let c = molar solubility, concentration of Ba
• Ksp = c x (2c)2 = 4c3
• 6 x 10-10 = 4c3
• c = 5.3 x 10-4

parts of atmosphere - 4 regions  

• troposphere - where people live, goes to about 12 km
• stratosphere - reaches 50 km
• mesosphere - reaches 85 km
• thermosphere - past mesosphere
• temperature change opposite in successive region
• pressure change remains constant (decreases as elevation increases)

atmospheric composition - made up of mostly nitrogen, oxygen  

• nitrogen = 78%
• oxygen = 21%
• argon, carbon dioxide, water vapor = next most commonly found gases
• oxygen reacts more than nitrogen (triple bond)
• photodissociation - photons from the sun breaking chemical bonds
  • O2 + energy >> 2O
  • mostly dissociated at very high elevations
• photoionization - loss of electron due to molecule absorbing radiation
• ozone (O3) forms from photodissociated oxygen and oxygen molecules
  • most ozone found in stratosphere
• chlorofluorocarbons (CFC) - depletes the ozone layer
  • harmless in lower atmosphere
  • reaches higher atmosphere >> more radiation >> more photodissociation
  • chloride w/ oxygen forms chlorine monoxide

troposphere composition - like atmosphere, mostly nitrogen/oxygen  

• acid rain - due to nitrogen oxides in air
  • NOx from fossil fuel burning
  • kills animals, corrodes metals/stone
• carbon monoxide - along w/ cyanide, bonds to hemoglobin better than oxygen
• photochemical smog - stagnant air mass due to nitric oxide
  • ozone acts as pollutant in lower atmosphere
• carbon dioxide/methane - contributes to greenhouse effect

seawater - contains 97.2% of world's water, aka saline water  

• salinity - grams of salts per 1 kg of seawater
• salt, bromine, magnesium commercially extracted from seas
• desalination - removing salts from seawater
  • distillation - makes the salts precipitate out
  • reverse osmosis - extra pressure >> solvent goes towards more dilute side

freshwater - amount of oxygen = quality of water  

• 9 ppm of oxygen in fully saturated water
• aerobic bacteria uses oxygen to oxidize biodegradable materials
  • excess biodegradable materials >> bacteria uses up more oxygen for oxidation
  • not enough oxygen >> aerobic bacteria dies >> anaerobic bacteria take over
• eutrophication - increasing amount of dead/decaying plant matter

green chemistry - using chemical products that preserve the environment  

• preventing waste > cleaning up waste
• industries need as little waste as possible, waste shouldn't be toxic
• chemical processes need to be energy-efficient
• raw materials should be renewable
• no auxiliary substances used (or at the very least, harmless auxiliary substances)
  • solvents not consumed in reaction but released into atmosphere

water purification - water chlorination produces trihalomethanes (THM)  

• several THMs are carcinogens

spontaneous reaction - occurs w/o any outside intervention

• usually exothermic, but can also be endothermic
  • ammonium salts dissolving in water is endothermic
• DE = q + w
  • q = heat absorbed by system from surroundings
  • w = work done on surroundings by system
• reversible process - can be restored to its original state using same path
  • results in no net change in system, surroundings
  • includes all chemical systems at equilibrium
• irreversible process - can only be restored using a different path
  • different q, w values from initial process
  • includes all spontaneous reactions
• can be fast or slow

Find the conditions surrounding H2O(l) >> H2O(s)  

• temperature decreases for water to turn to ice
• energy leaves water
  • DH is negative
• product is more orderly/structured than reactant
  • DS is negative

entropy - the amount of disorder in a system

• DS = Sfinal - Sinitial
• positive value >> more disorder
• negative value >> less order
• 2nd law of thermodynamics - entropy increases for all spontaneous reactions
  • S = 0 for reversible processes
• 3rd law of thermodynamics - entropy = 0 at absolute zero (0K)
  • translational motion - entire molecule moving
  • vibrational motion - periodical mov’t to/away
  • rotational motion - spinning mov’t
  • increase temperature >> more mov’t >> more entropy
  • decrease temperature >> less mov’t >> less entropy
• Ssolid < Sliquid < Sgas
• å S = å Sproducts - å Sreactants

Arrange the following changes from least increase in entropy to greatest increase in entropy: 1g of ice warmed by 1°, 1g of ice melted, 1g of water frozen, 1g of water evaporated.  

• water to liquid is decrease in entropy
• phase changes increase in entropy more than just increasing temperature
• gas has more entropy than liquid
• 1g of water frozen, 1g of ice warmed by 1°, 1g of ice melted, 1g of water evaporated

Gibbs free energy - free energy

• G < 0 >> spontaneous reaction
• G = 0 >> reaction at equilibrium
• G > 0 >> nonspontaneous reaction (reverse reaction could be spontaneous)
• G = H - TS
• DG =DH - T (DS)
• always negative when H negative, S positive
• always positive when H positive, S negative

Find the conditions surrounding the stretching of a rubber band.  

• needs energy to stretch rubber band
  • DG is positive (nonspontaneous event)
• rubber band becomes more orderly when stretched
  • DS is negative
• release of heat when rubber band is stretched
  • DH is negative

Calculate DG (in kJ) for Mg(s) + 1/2O2 >> MgO(s) at 300K  

• Given:
  • DHf of MgO = -601.8
  • DS° of Mg(s) = 32.7 J
  • DS° of O2 = 205 J
  • DS° of MgO(s) = 26.9 J
• DS = 26.9 - (1/2 x 205 + 32.7) = -108.3 J = -0.1083 kJ
• DG = -601.8 - 300(-0.1083) = -569.31 kJ

nonstandard free energy - value differs from standard value at different conditions

• DG = G° + RT lnQ
  • R = constant (8.314 J/mol-K)
  • Q = equilibrium constant
• at equilibrium,G° = -RT lnKeq
• G < 0 >> Keq > 1
• G = 0 >> Keq = 1
• G > 0 >> Keq< 1

Calculate the Keq for 2NO2 (g) >> N2O4 (g) at 25°C  

• Given:
  • DHf of NO2 = 33.84 kJ
  • DHf of N2O4 = 9.66 kJ
  • DS° of NO2 = 240.4 J
  • DS° of N2O4 = 304.3 J
  • R = 8.314
• DS = 304.3 - 2(240.4) = -176.5 J = -0.1765 kJ
• DH = 9.66 - 2(33.84) = -58.02 kJ
• DG = -58.02 - (273+25)( -0.1765) = -5.423 kJ = -5423 J
• DG = -RT lnKeq
• -5423 = -(8.314)(298) lnKeq
• lnKeq = 2.1888
• Keq = e2.1888 = 8.92

oxidation-reduction reactions - oxidation states change

• oxidation - increasing oxidation number, losing electrons
• reduction - decreasing oxidation number, gaining electrons
• reduction always accompanies oxidation (and vice versa)
• oxidizing agent (oxidant) - makes it possible for another substance to get oxidized
  • gets reduced
• reducing agent (reductant) - makes it possible for another substance to get reduced
  • gets oxidized
• half-reactions - show either oxidation or reduction alone
  • electrons cancel each other out when combined into single equation

For the following equations, determine the oxidizing agent and reducing agent.  

• remember, oxidizing agent gets reduced, reducing agent gets oxidized
• I2O5 + 5CO >> I2 + 5CO2
  • I+5 >> I0, gets reduced - oxidizing agent
  • C+2 >> C+4, gets oxidized - reducing agent
• H2O2 + C2H4 >> H2O + C2H4O
  • C-2 >> C-1, gets oxidized - reducing agent
  • O-1 >> O-2, gets reduced - oxidizing agent

steps for balancing a redox reaction

• divide equation into 2 half-reactions
• balance each half-reaction
  • balance all elements other than H and O
  • balance O atoms by adding H2O
  • balance H atoms by adding H+
• multiply each half-reaction by integer so that # of electrons lost equals # of electrons gained
  • add together half-reactions, cancel out all spectator substances
• check
• if in basic solution, add OH- to both sides in order to cancel H+ ions

Balancing MnO4-1 + Fe+2 >> Mn+2 + Fe+3  

• Mn+7 >> Mn+2
  • MnO4-1 >> Mn+2 + 4H2O
  • MnO4-1 + 8H+ >> Mn+2 + 4H2O
  • oxidation number of Mn -5
• Fe+2 >> Fe+3
  • no need to balance
  • oxidation number of Fe +1
• 5Fe+2 + MnO4-1 + 8H+ >> 5Fe+3 + Mn+2 + 4H2O

voltaic (galvanic) cells - electrons transferred through external pathway instead of directly

• oxidation occurs at anode
• reduction occurs at cathode
• half-cells must stay neutral >> cations migrate over to cathode, anions migrate over to anode across salt bridge (won’t react w/ other ions)
• electrons always flow towards anode

cell EMF (electromotive force) - aka cell potential/voltage

• Ecell = Ecathode - Eanode
  • positive value >> spontaneous
• standard hydrogen electrode (SHE) - has reduction potential of 0V
  • used to measure voltage of half-reactions
• intensive property >> changing coefficients in reaction won’t change value
• E more positive >> greater tendency to reduce (at cathode)
  • strongest oxidizing agent >> most easily reduced
  • strongest reducing agent >> most easily oxidized
• G° = -nF(E°)

For the following batteries, determine the anode/cathode and voltage  

• remember, more positive reduction potential means cathode, more negative potential means anode
• Li-Cd battery
  • Li reduction potential = -3.05 (anode)
  • Cd reduction potential = -0.40 (cathode)
  • voltage = -0.40 - (-3.05) = 2.65
• Al-Cd battery
  • Al reduction potential = -1.66
  • Cd reduction potential = -0.40
  • voltage = -0.40 - (-1.66) = 1.26

effect of concentration on cell EMF - depends on Nernst equation

• E = E° - RT / nF lnQ
  • = E° - 2.303RT / nF logQ
  • = E° - 0.0592 V / n logQ at 298K
• reactant concentration increase >> emf increase
• product concentration increase >> emf decrease

concentration cell - cell based on emf generated from difference in concentration

• uses same substance
• will operate until concentration equal
• use Nernst equation to figure out
• emf of voltaic cell decreases as it discharges

battery - portable, self-contained electrochemical power source w/ 1 or more voltaic cells

• use multiple voltaic cells >> greater voltage
• primary cells - can’t be recharged
• secondary cells - can be recharged from external power source
• lead-acid battery - 2V battery w/ lead dioxide as cathode and lead anode
  • 6 strung together in 12-V automotive battery
  • can be recharged
• alkaline battery - most common primary battery
  • manganese oxide and graphite mixed in cathode, zinc anode
  • emf of 1.55
• nickel-cadmium (nicad) battery - most common rechargeable battery
  • environmental hazard, increases weight of batteries
• nickel-metal-hydride (NiMH) battery - uses alloy for anode
• litium-ion (Li-ion) battery - has higher energy density than nickel-based batteries
• fuel cells - uses conventional fuels, not batteries (not self-contained)
  • most promising system uses hydrogen/oxygen, forms water as only product

corrosion - metal converted to unwanted compound due to environment

• rusting - forms Fe2O3 * xH2O from iron/oxygen
  • rust usually deposits at cathode (largest supply of oxygen)
• paint/metal coatings added to protect against corrosion
• galvanized iron - zinc layer added on to iron
  • zinc gets corroded before iron
  • sacrificial anode - oxidized first to protect another cathode

electrolysis - nonspontaneous redox reactions started by outside energy source

• electrolytic cell made of 2 electrodes in molten salt or solution
• electrolysis of molten salts needs high temperatures
• inert electrodes - serve as surface where oxidation/reduction occur
• active electrodes - participate in oxidation/reduction process
• electroplating - uses electrolysis to deposit thin metallic layer on another metal

radioactivity - tendency to decay, emit particles

• radionuclides - radioactive nuclei
• radioisotopes - atoms w/ radioactive nuclei
• alpha radiation - emitting alpha particles (helium-4 nuclei)
  • atomic number decreases by 2
  • atomic mass decreases by 4
• beta radiation - emits electrons from unstable nucleus
  • neutron >> proton
  • atomic number increases by 1
  • atomic mass stays the same
• gamma radiation - emits high-energy photons
  • accompanies other radioactive emissions
  • doesn’t change atomic number/mass
• electron capture - nucleus gains electron from surrounding electron cloud
  • proton >> neutron
  • atomic number decreases by 1
  • atomic mass stays the same

nuclear stability - depends on multiple factors

• neutron-proton ratio - neutron # > proton # to ensure stability in atoms w/ large atomic numbers
  • belt of stability - area in which all stable nuclei are found
  • all nuclei w/ 84+ protons are radioactive
  • nuclei above belt of stability use beta emission >> increases proton #, decreases neutron #
  • nuclei below belt of stability use electron capture, positron emission >> decrease proton #, increase neutron #
  • nuclei with over 84 protons use alpha emission
• radioactive series - aka nuclear disintegration series
  • 3 naturally occurring series end with lead
  • stability not achieved w/ a single emission (needs a series of emissions)
• magic numbers - 2, 8, 20, 28, 50, 82 protons; 2, 8, 20, 28, 50, 82, 126 neutrons
  • usually more stable when w/ these numbers

nuclear transmutations - changing the atom’s identity  

• 1st conversion done by Rutherford
• nuclear reactions can start by striking nuclei w/ particles
• particle accelerators - slams charged particles against nuclei
  • needs high speed to overcome electrostatic repulsion between particles and nuclei
  • not used for neutrons (not repelled by nuclei)
• transuranium elements - elements w/ atomic number greater than 92

Xenon-118 undergoes electron capture to become what?  

• atomic number decreases by 1
• atomic mass stays the same
• 118Xe + -1e0 >> 118I

Carbon-14 undergoes betay decay to produce what?  

• atomic number increases by 1
• atomic mass stays the same
• 14C >> 14N + -1e0

radioactive decay rates - 1st-order kinetic process

• half-life - time required for 1/2 of given quantity to react
• k = 0.693 / t1/2
• ln (Nt / N0) = -kt

An old piece of wood has carbon-14 activity of 11.7 disintegrations per minute per gram. Current carbon-14 activity is 15.3 disintegrations per minute per gram. The half-life of carbon-14 is 5714 years. How old is that old piece of wood?  

• Given:
  • t1/2 = 5714
  • Nt = 11.7
  • N0 = 15.3
• k = 0.693 / 5714 = 1.2 x 10-4
• ln (11.7 / 15.3) = -(1.2 x 10-4)t
• t = 2212 years old

radioactivity detection - radioactivity found everywhere

• Geiger counter - creates pulse whenever radiation enters tube
  • radiation causes matter to ionize, create electrical current
• scintillation counter - measures radiation based on light produced when radiation hits a phosphor
  • phosphor - substance that gives off light when hit by radiation
• radiotracer - radioisotope used to trace the path of an element through a reaction

nuclear fission - fragmenting heavy nuclei to generate energy

• chain reactions - 1 leading to another, multiplying number of reactions
• critical mass - amount of fissionable material needed to maintain chain reaction
  • supercritical mass - in excess of critical mass
• control rods - absorbs neutrons to regulate chain reaction
  • keeps reaction self-sustained, prevents overheating
• moderator - slows down neutrons
• cooling liquid - carries off heat generated by fission (can also work as the moderator)
• problems w/ storing the radioactive wastes from fission

nuclear fusion - combining/fusing light nuclei

• used by sun to produce energy
• doesn’t produce any radioactive substances
• requires at least 40,000,000 K to start reactions
• tokamak - uses strong magnetic fields to contain/heat reaction
  • only reached temperatures of 3,000,000K

biological effects of radiation - radiation >> excitation or ionization of matter

• excitation >> electrons go to higher energy states, moves more
• ionization >> electron gets removed from atom
  • creates free radicals (has 1+ unpaired electrons)
  • can attack other compounds
• gray (Gy) - absorption of 1J of energy per kg of tissue
  • 100 rads = 1 Gy
• radon - radioactive noble gas fromed from uranium-238
  • no direct chemical effects when inhaled
  • short half-life >> can produce radiation very quickly

Please click below to download the AP Chemistry outline for 'Chapter 1 - Chemical Foundations', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for, 'Chapter 2 - Atoms, Molecules, and Ions', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for, 'Chapter 3 - Stoichiometry', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for, 'Chapter 4 - Types of Chemical Reactions and Solution Chemistry, from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for, 'Chapter 5 - Gases', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for, 'Chapter 6 - Thermochemistry', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for, 'Chapter 7 - Atomic Structure and Periodicity', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for 'Chapter 8 - Bonding: General Concepts', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for 'Chapter 9 - Covalent Bonding: Orbitals', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for 'Chapter 10 - Liquids and Solids', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for 'Chapter 11 - Properties of Solutions', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

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Please click below to download the AP Chemistry outline for 'Chapter 13 - Chemical Equilibrium', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for 'Chapter 14 - Acids and Bases', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

Please click below to download the AP Chemistry outline for 'Chapter 15 - Applications of Aqueous Equilibria', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

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Please click below to download the AP Chemistry outline for 'Chapter 21 - The Nucleus: A Chemist's View', from the Zumdahl's Chemistry, 5th Edition Textbook. These AP Chemistry notes will cover the key topics discussed in this chapter.

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Please click below to download the AP Chemistry unit notes for the Holt’s Modern Chemistry Textbook. These AP Chemistry unit notes will cover the key topics discussed in this unit.

Please click below to download the AP Chemistry unit notes for the Holt’s Modern Chemistry Textbook. These AP Chemistry unit notes will cover the key topics discussed in this unit.

Please click below to download the AP Chemistry unit notes for the Holt’s Modern Chemistry Textbook. These AP Chemistry unit notes will cover the key topics discussed in this unit.

Please click below to download the AP Chemistry unit notes for the Holt’s Modern Chemistry Textbook. These AP Chemistry unit notes will cover the key topics discussed in this unit.

Please click below to download the AP Chemistry unit notes for the Holt’s Modern Chemistry Textbook. These AP Chemistry unit notes will cover the key topics discussed in this unit.

Please click below to download the AP Chemistry unit notes for the Holt’s Modern Chemistry Textbook. These AP Chemistry unit notes will cover the key topics discussed in this unit.

Please click below to download the AP Chemistry unit notes for the Holt’s Modern Chemistry Textbook. These AP Chemistry unit notes will cover the key topics discussed in this unit.